Smith 1
Chemistry Review Unit I: Matter and Energy The Periodic Periodi c Table Table Groups of the Period Table
• Group 1a: Alkali Metals- Lithium, Sodium, Potassium, Rubidium, Cesium, Francium • Group 2a: Alkaline Earth Metals- Beryllium, Magnesium, Calcium, Strontium. Barium, Radium • Group 3-12: Transition Metals • Group 13: Boron Family- Boron, Aluminum, Gallium, Indium, Thallium • Group 14: Carbon Family- Carbon, Silicon, Germanium, Tin, Lead Antimony, Bismuth • Group 15: Nitrogen Family- Nitrogen, Phosphorous, Arsenic, Antimony, • Group 16: Oxygen Family- Oxygen, Sulfur, Selenium, Tellurium, Polonium • Group 17: Fluorine Family/Halogens- Fluorine, Chlorine, Bromine, Iodine, Astatine • Group 18: Noble Gases/Inert Gas Group- Helium, Neon, Argon, Krypton, Xenon, Radon Classifying The Elements
I. Demitri Mendeleev (1869)- Russian Scientist • Concluded that physical and chemical properties of elements appear in regular intervals when listed by increasing atomic mass. • Not current classification II. Henry Mosely (1913)- English Scientist • Used x-rays to identify atomic mass. • He concluded that physical & chemical properties of elements were listed by increasing atomic number. • Modern classification system Basic Units of Measurement
• Length- Measure distance between two points a. Basic unit: Meter. Meter. Instrument: Meterstick
Smith 2 b. 1 km= 1,000m 1m=100cm 1cm= 10 mm • Mass- Measures the quantity of matter a. Basic unit: Gram. Instrument: Triple Beam Balance b. 1 kg=1,000g 1g=1,000 mg b y an object • Volume- Measures amount of space occupied by a. Basic unit: Liter Instrument: Graduated Cylinder b. 1L=1,000mL 1mL= 1cm^3 c. Volume is represented by Length, Width and Height (V=LWH, (V=LWH, for Regular Shape) d. Vo=Vwto-Vw -Volume for irregular solid; Vw= Initial level of water, Vwto- Final level of water • Temperature- Measures "hotness" or "coldness: a. Basic Unit- Degrees Celsius with Thermometer b. Freezing point H2O- 0C c. Boiling Point H2O- 100C d. International Scale (SI) Temperature- Measured with Kelvin Scale (Based off of absolute zero) - 0C= 273K, 0K=-273C • Time - Measures duration a. Basic Unit: Second with Clock b. 60 Sec= 1 min 60 min=1 hour Significant Figures
A. All numbers that are actually read plus one estimated guess. I. Initial zeros are never significant. 0.0203: Only 3 significant figures II. All whole numbers are always significant. 2039- 4 significant figures II. Zeros are significant if they are between who le numbers. 2004- 4 significant figures IV. IV. Zero is significant if it's to the right of the wh ole number and to the right of the decimal point. 0.09036- 4 significant figures V. Final zeros are ambiguous a mbiguous
Smith 3 VI. For addition and subtraction, the answer can contain as many decimal places as the least accurate value. VII. For multiplication and division, the answer can only contain as many significant digits as the least accurate value. What is Chemistry? I. It is the study of:
a nd properties of matter. • The composition, structure and • The changes which matter undergoes. • And the energy which occupies these changes. II. Matter
• Anything that has mass and volume • Divided into two classes: Mixtures and Substances III. Substances • All homogeneous matter a. Fixed composition b. Always made of the same amount of matter Types of Substances: Elements and Compounds Compound s • 2 Types I. Elements • Contain all atoms of a single type • Cannot be decomposed • Common elements: Oxygen, Nitrogen, Carbon, Hydrogen 2. Compounds • Contains more than 1 type of element • Chemically combined in fixed ratio • Can be decomposed chemically but not physically • Binary Compounds- 2 Elements, like NaCl • Ternary Compounds- 3 Elements, like KNO3
Smith 4 III. Mixtures
• Consists of 2 more elements that differ in property and composition • Substances are physically mixed • The composition/ratio of substances vary • Can be separated physically • 2 Types: Homogenous and Heterogenous a. Homogenous- Uniformly mixed throughout the mixture aa. Also called Solutions i. Dissolved particles in solution (Na in H2O) ii. Aqueous solution dissolved in water bb. Distillation - A method used to separate parts of a homogenous mixture based on their boiling points. b. Heterogenous- Multiple components that are randomly distributed. aa. Filtration is a method of using a filter to physically separate the mixture i. Material collected ia s filtrate, material left behind is the residue. bb. Chromatography- Separates part of the mixture physically as they have a different rate of moving up filter paper.
Properties Physical Properties
I. A quality of a substance that can be observed or measured without changing the substance's composition. II. Examples: Color, solubility, odor, hardness, density, melting point, boiling point, luster (senses). Chemical Properties
I. The ability of a substance to undergo unde rgo a chemical reaction & to form a new substance II. A substance must undergo a chemical change to observe a chemical property III. Examples: Rust, burn, rot, decompose, ferment, explode, corrode.
Smith 5 Metallic Properties
I. Luster, good conductors of heat & electricity, malleable (hammer into shapes), ductile (ability to be drawn into wires), hard. Non-Metallic Properties
I. Dull, brittle, soft (if solid), poor conductors of heat & electricity, no free electrons. Metalloids/Semi-Metal Properties
I. Properties of both metals & non metals II. Shiny, but brittle, like Silicon Physical vs. Chemical Changes
I. A physical changes in form, but does not become something new II. Dissolving, Melting, Boiling, Freezing, Cutting III. A chemical change forms a new substance, energy always accompanies a chemical changes Matter-Continued
I. Solids • Matter is arranged in a regular, rigid pattern • Definite shape and volume • Crystalline structure- crystals arranged in a repeated geometric pattern (Like ice, strong intermolecular attraction) II. Liquids • Particles not held as tightly together • Able to move past one another (flow) • Definite volume, no definite shape III. Gases • Minimal attractive holding particles together • No definite shape or volume (takes shape and volume of container) IV. Psuedosolids • Lacks crystalline structure
Smith 6 • Supercooled liquids-molecules move over one another in time • Like glass and some plastics Phase Changes
• All phase changes are accompanied with either a loss or gain of energy • An element, compound or mixture can exist as a solid, liquid or gas I. Endothermic Reactions • Phase change that requires the gain of heat • Melting/Fusion- Solid becomes liquid • Evaporation/Boiling/Vaporization- Liquid becomes gas • Sublimation- Solid turns into gas directly (Substances that sublime have high vapor pressure and low intermolecular forces of attraction) II. Exothermic Reactions • All of these changes require the loss or release of energy/heat • Freezing/Solidification-Liquid becomes solid • Condensation- Gaseous substance becomes liquid • Deposition- Gaseous substance turns directly into a solid Main Types of Energy
I. Kinetic (Ke)- Energy of Motion II. Potential (Pe) Stored Energy III. Chemical Energy (Ce)- Energy associated with a chemical change A.Ke=1/2mv^2 B.Pe=ugh (mass x distance from the ground x gravity) IV.Heat Energy A.Amount of energy transferred from one substance to another B.Can be measured using a calorimeter C. Calories (cal) or Joules (J) measure heat gain or loss
Smith 7 D.To convert from cal or J to Cal or KJ, divide by zero Law of Conservation of Energy • Energy is neither created nor destroyed • Energy can be transferred from one substance to another • Or energy can be transferred into a new form of energy • The total abound of energy will remain the same • Example: Gas burns in engine (Chemical to heat), Car moves (Mechanical to kinetic) Thermometry
• Temperature- The measure of the average kinetic energy of the particles of a substance • Heat- Flows spontaneously from a hot body to a cold body • Body heat to chair, boiling water to hand, burned hand to icepack • Temperature Scales • Degree Celsius- Most commonly used, 2 fixed po ints (0 which is melting/freezing of water, and 100, which is boiling/condensation point of water) Values increase by 1 • Kelvin- Contains theoretically the lowest possible temperature, has never been exactly reached, absence of all kinetic energy • K=C+273, C=K-273 Measurement of Heat Energy
• The amount of heat given off or absorbed in a reaction can be calculated by:
Q=MC!T --Q=Heat (Joules or calories) --M=Mass of substance °
--C=Specific heat capacity of substance (J/G C) --!T=Temperature Final-Temperature Initial • Q=MC!T is used only when there is a change in temperature
Smith 8 °
• Specific Heat Capacity- Amount of heat needed to raise 1g of a substance by 1 C (Water=4.18 J/G C) °
Heat of Fusion
• The amount of heat needed to melt 1g of a substance • Q=MHf is used when calculating how much heat is absorbed when a substance melts • Remember: • Heat absorbed during melting goes into raising the potential energy of the Substance • Kinetic energy is constant (constant temperature therefore you cannot use Q=MC!T • The value for heat of fusion is 334 J/G Heat of Vaporization
• The amount of heat needed to vaporize/boil 1g of a substance • Q=MHv is used to calculate how much energy is absorbed when a substance vaporizes • The heat of vaporization of water is 2260 J/G Mole Concept-Avogadro’s Number
• Based off of the atomic mass of Carbon • 1 Gram of Hydrogen=1 Mole/ 6.02x10^23 amu/particles • 12 Grams of Carbon= 1 Mole/ 6.02x10^23 amu/particles • 24 Grams of Magnesium= 1 Mole/ 6.02x10^23 amu/particles • Example: Magnesium Nitrate Mg(NO3)2 • Mg (1)(24)=24g • N (2)(14)=28g
Mg(NO3)2 = 148g = 6.02x1023 Particles
• O (6)(16)=96g
24+28+96=148
• Atoms and molecules are too small to count, so we count them in liege quantities • The number of atoms of carbon present in 12 grams • The mass of one mole of a substance can be found by determining its gram-formula mass • To convert grams to moles, use formula found in reference table
Smith 9 • Given Mass/Gram-Formula Mass Gas-Mole Concept
• 1 Mole of any gas at STP= 22.4L • 1 Mole of H2(g) =22.4 L, 1 mole of Cl2(g)=22.4L • Diatomic molecules that exist in nature= • I2, Br 2, Cl2, F2, O2, N2, H2 • I BRought CLay For Our New House • Monoatomic Molecules - Noble Gases- He, Ne, Ar, Kr, Xe • To solve moles for gases- Moles=Given Liters/Liters Per Mole (22.4 L)
Smith 10 Percent Composition
• The percent by mass of each element in the compound. • The percent composition of a compound consists of a percent value for each different element in the compound • K 2CrO4 • K= 40.3% • Cr=26.8% • O=32.9% • The percents MUST total 100% • The percent by mass of an element in a compound is the number of grams of the element divided by the mass in grams of the compound, multiplied by 100% I. K (39)(2)=78g. Cr (1)(52)=52g. O (4)16) = 64 - Configure Gram-Mass Formula • 78+52+64=194g II. Use proportion 78/194 = 40.3% 52/194=26.8% 64/194= 32.9%
Gas Laws, Matter Kinetic Molecular Theory for Ideal Gases
• Studies of gas behavior have led to a model referred to as the ideal gas model-based off of several assumptions • A gas is composed of individual particles which are in a con stant, straight line motion • Gas particles are separated great distances relative to their size. The volume of gas particles are not considered. • Gas participles are considered as having no attraction to each other. • The collision theory states that a reaction is most likely to occur if the reactant particles collide with proper energy and orientation (Sufficient amount of energy and p roper angle & geometry) Deviations from the Gas Laws
• The ideal gas model does not exactly represent real gases under all conditions • Hydrogen and Helium are the two most ideal gases- no real gas follows the ideal model under all conditions of temperature and pressure
Smith 11 • Deviations occur because model is not perfect • This is because gas particles have volume and exert some attraction for each other • These factors because significant under conditions of high temperature and low pressure and decreased velocity due to increased molecular mass • Conditions of high temperature and low pressure are ideal Gas
• The space between molecules in a gaseous phase is about 1,000 times greater than in liquid or solid phase. • Molecules possess greater kinetic energy and have o vercome the attractive forces that hold them together. • The density of a gas is lower as compared to solid or liquid • *In the gas phase, molecules vibrate, rotate and translate. This allows them to fill the volume of the container in which they are held • High Temperature/Lower Volume= Increased collision, smaller particles move faster. • High Temperature/Low Pressure= ideal • High Pressure/Low Temperature=real Gas Laws
I. Boyle’s Law • Temperature=Constant • P1V1=P2V2 • V2=P1V1/P2
PV=Constant
• As Volume increases, Pressure decreases (Inverse Proportions) II. Charles’ Law • Pressure=Constant • V1/T1=V2/T2 • V2=V1/T1 • As temperature increases, volume increases • Directly proportional
Smith 12 III. Gay Lussac’s Law • Volume=Constant • P1/T1=P2/T2 • P1T2=P2T1 • P2=P1T2/T1 • As pressure increases, temperature increases • Directly proportional
Gas Laws for Closed Systems
• P1V1/T1=P2V2/T2 - Charles Law - Temperature always in Kevin, increased volume in an expandable container (constant pressure) • P1V1/T1=P2V2/T2 Gay-Lussac’s Law - Temperature always in Kelvin, increased pressure in a rigid container (constant volume) • Increase in average kinetic energy causes increase in momentum, which causes an increase in collision frequency • P1V1/T1=P2V2/T2 - Boyle’s Law- Temperature always in Kelvin, Increase in pressure causes a decrease in pressure (temperature remains constant) • Increase in external pressure causes increase in external collisions, which causes an expandable container pushed into smaller container Phases in Detail-Gases
• Molecules spread out and fill spaces- they are given due to their weak intermolecular forces of attraction • There are large spaces between gas particles • Evaporation • Takes place at all temperatures on the liquid/vapor boundary • Vapor-A gaseous phase of a substance that is a liquid or solid at normal conditions • Once liquid particles have absorbed enough energy to overcome attractive forces, they become vapor
Smith 13 • Vapor Pressure- Gas particles exert pressure on the liquid when they evaporate • In a closed system, the pressure increases • Evaporation increases with an increase in temperature • 1 ATM=760 Torr=760 mm Hg, are also values of standard pressure • Normal Boiling Point= When the vapor pressure=atmosphere pressure. • When a substance boils, evaporation occurs throughout the liquid • Also measures the strength of intermolecular forces • *If Vapor Pressure is high, attraction between molecules is weak • *If Vapor Pressure is low, attraction between is strong
History of the Atom
I. John Dalton’s Atomic Theory (1803) • All matter is made up of tiny, indestructible particles called atoms • All atoms of a given element have identical physical & chemical properties • Atoms are neither created nor destroyed (Law of conservation of mass) • Atoms of different elements form compounds in while number ratios • Some of these postulates now have exceptions: • Atoms can be broken apart in nuclear reactions • Atoms of a given element can have different physical and chemical properties (isotopes) II. J.J. Thomson’s Plum-Pudding Model (1897) • His model portrays the atom as a big ball of positive charge that contains small particles of negative charge embedded in it. • Discovered the charge of an electron by observing cathode rays in a cathode ray tube • From his observation, he concluded that cathode rays are streams of negatively charged particles with mass • Another scientist, Millikan, was able to determine the mass of an electron based upon Thomson’s work (1909). III. Rutherford’s Model (1909) • Made two key observations based on his “gold foil” experiment
Smith 14 • He disproved Thomson’s model • He bombarded a thin piece of gold foil with positively charged alpha (positive cha rge) particles (much smaller than the atom) • Proved the nucleus to be positive • Observations: I.
Almost all the alpha particles passed through foil without deflection
II.
Small percent slightly deflected
III.
Some were largely deflected
IV.
A few even reflected back in the direction from where they had came
• Conclusion: (1911): Atom is mostly empty space and all of the positive charge in an atom is concentrated in a small, dense core (nucleus). • This area is positive, since positively charged particles were deflected from it (repelled) • Atomic Mass= Sum of protons and neutrons • Atomic Number= Number of protons • Number of protons=Number of electrons • Last level= Valence electrons IV. Bohr Planetary Model (1913) • Model displayed electrons traveling in orbits around the nucleus • Electrons are only found in orbitals (principle energy levels) not in between • The principle energy levels (PEL) approximates how far the electrons are from the PEL uncles PEL Shell Max # Of Electrons 2(N)2 N=PEL 1
K
2
2
L
8
3
M
18
4
N
32
5
O
50
6
P
72
7
Q
98
• The electrons’ distance from the nucleus is related to their specific amount of energy (quanta)
Smith 15 • As you move away from the nucleus, the energy in each PEL Increases. • Like climbing stairs, further you go=more energy • Ground State-When electrons are in lowest energy level • Quantum Leap- When electrons jump between energy levels • 2 electrons can only absorb a fixed amount of energy (quanta) to move to higher level • Electrons can only jump to levels that aren’t completely filled with electrons • Heat, light and electricity are all stimuli that can excite an electron • Excited State- Electrons are in higher energy levels. Acquired when an electron absorbs energy and becomes unstable. Electrons quickly return to ground state, emitting the same amount of energy absorbed, usually in some form of light. • Every element gives off a unique pattern of colors (line spectrum) which can be used to identify the element • Planck’ s constant, h = 6.63 ¥ 10-34 J s
Electron Configuration
• An electron configuration tells you how many electrons there are in each energy level • 1 Mg (2-8-2) has 2 electrons in PEL 1, and 8 electrons in PEL 2 • The amount of numbers in each electron configuration tells you how many electron levels are occupied with electrons • 2n2 - energy levels for max (for max) • 2-8-2- Ground State, 2-7-3- Excited State • Last=Valence Electrons Valence Electrons
• The electrons in the outermost energy level of an atom (last # in the electron configuration) • 2-8-3 has 3 valence electrons • Valence electrons can determine the chemical properties of an element The Kernel
• Includes the nucleus and all non-valence electrons Quantum Numbers
• Schrödinger- Mathematically treated the electron as a wave • The 4 quantum numbers in Schrödinger’s equation are used in describing electron behavior
Smith 16 • N,L,M,S • Principle Quantum Number- N • Second quantum number indicated by L describes sublevels. Each energy level (N) has n sublevels. • Level
Sublevel
1
1 s
2
1,2 s,p
3
1,2,3 s,p,d
4
1,2,3,4 s,p,d,f
• The third quantum number m represents the number of orbitals in a sublevel • Sublevels- only 1 orbital m=0 • Sublevels- 3 orbitals (x,y,z) m=0±1 • Sublevels- 5 orbitals m=0±1±2 • Sublevels 7 orbitals m=0±1±2 • Only 2e- in each orbital • Spin Quantum Number- Describes the spin of an atom (Pauli) clockwise • Pauli’s Exclusion Principle- No two electrons in an atom can have the same set of 4 quantum numbers • Examples: I. Hydrogen- 1s1 II. He 1s2 III. Lithium - 1s22s1 IV. Carbon 1s22s22s2 V. Neon 1s22s22p6 VI.Magnesium - 1s22s22p63s2
Smith 17 Bonding Ionic Bonding
• Compounds composed of cations and anions are called ionic compounds • Characterized by the transfer of electrons- Representative unit is the formula unit • Composed of metal cations and nonmetal anions. • Cation- Ion with positive charge • Anion- Ion with negative charge • Anions and cations have opposite charges and attract one another with electrostatic forces • Properties: I.
Solid at room temperature
II.
Have high melting points
III.
Conduct an electric current when Dissolved/melted in water
• Use brackets diagram to illustrate electron transfer based off of oxidation numbers and Valence electrons - to satisfy the octet rule (to obtain 8 valence electrons to achieve stability-like the noble gases) • Example: Calcium and Chlorine (Metal and nonmetal) Calcium’s electron configuration: 2-8-8-2 Chlorine’s electron configuration: 2-8-7 You need to remove 2 electrons from calcium to achieve 8 valence electrons *Use Lewis Electron Dot Diagram* - Depicts valence electrons And 2 electrons must be given to chlorine (Cl is diatomic), so one electron goes to each chlorine.
Smith 18
Covalent Bonding
• Characterized by the share of electrons (like a tug-of-war between elements) to achieve electron configuration of noble gases • Representative unit is a molecule • Nonpolar and Polar • Nonpolar: Bonding electrons are shared equally (Like N2, O2, Cl2, H2) • Polar: Bonding electrons shared unequally • Polar Molecule*: Asymmetrical molecule, one side is more negative than the other is positive. (AKA, a dipole) • Nonpolar Molecule*: Symmetrical molecule, charges are balanced • The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge • Use difference of electronegativity to determine most probably type o f bond Electronegativity Difference
Most Probable Bond
Example
0.0-0.4
Nonpolar Covalent
H-H (H2) (0.0)
0.4-1.0
Moderately Covalent
H-CL (HCl) (0.9)
1.0-2.0
Very Polar Covalent
H-F (HF) (1.9)
"2.0
Ionic
Na+Cl- (2.1)
• Electronegativity: The ability of an atom to attract electrons when the atom is in a compound • Single Covalent Bond: Bond formed when when two atoms share a pair of electrons (Like H:H)-Depicts the sharing of two electrons • Double Covalent Bond: A bond in which two atoms hare two pairs of electrons (like O::O), Oxygen has 6 valence electrons, and needs 8 to follow octet rule, so oxygen shares two with oxygen) • Triple Covalent Bond: A bond formed by sharing three pairs of electrons (like N:::N), Nitrogen has 5 valence electrons, and needs 3 more to follow octet rule- so, Nitrogen must share 3 with Nitrogen)
Smith 19 • Network Solids/Crystals: Solids in which all of the atoms are covalently bonded to each otherVERY high melting point. Examples are Diamonds, Silicon Carbide, Silicon Dioxide • Properties: I. Low melting/boiling points II.Tend to be soft, tend to be liquids, gases or soft solids III. Poor conductors of heat and electricity IV.Are molecules V. Are brittle VI. Nonmetallic • Coordinate Covalent Bonds: A covalent bond in which one atom contributes both bonding electrons. (Can depict in structural formula by drawing an arrow that points from the atom donating the pair of electrons to the atom receiving them) (like CO)
Metallic Bonding
• Can be described as a sea of electrons • The valence electrons are mobile and can drift freely from one part of the metal to another • Metallic bonds consist of the attraction of the free-floating valence electrons for the positively charged metal ions. These bonds are the forces of attraction that hold metals together (Cu would be considered a metallic bond) • Sea of electrons explains physical properties of metals: • Excellent conductors of heat and electricity • Malleable (Can be hammered and shaped) • Ductile (Can be made into wires) • Metal atoms are arranged in very compact and orderly patterns
Smith 20 Hydrogen Bonding
• Attractive forces in which a hydrogen covalently bonded to a very electronegative element (Fluorine is most electronegative, oxygen is second most electronegative) is also weakly bonded to an unshared electron pair of an electronegative atom. • Like OH, HF, NH (Hydrogen with Oxygen, Fluorine or Nitrogen) • For a Hydrogen bond to form, there must be a covalent bond present • Strongest of intermolecular forces • Extremely important of determining the properties of water and biological molecules, such as proteins Van der Waals Forces
• The two weakest attractions between molecules- named after Dutch chemist Johannes van der Waals (1837-1923). • Van der Waals forces consist of dipole interactions and dispersion forces • Dipole interactions occur when polar molecules a re attracted to one another-The electrical attraction involved occurs between the oppositely charged region of polar molecules (Like NaCl(Aq), Na+ will attach to O- and Cl- will attach to H+ (negative goes to positive, positive goes to negative) • Dispersion forces are the weakest of all molecular interactions and are caused by the motion of electrons. • Caused by the electron motion on one molecule affecting the electron motion on the other through electrical forces (electrons are negative..when an electron moves, it will repel another electron) Trends in Periodic Table I. Atomic Radii
• Trend in Period • Decrease left to right • Increase # of protons, increase attraction. For valence electrons= smaller radius • Trend in Group • Increase Top to bottom
Smith 21 • Inner electrons shield valence electrons • Reduces attractive forces= Bigger radius II. Ionic Radius
• Metals lose electrons • Radius decreases in size • Metals have larger radius than its ion • Nonmetals gain electrons • Radius increases in size • Nonmetals have smaller radius than its ion III. Electronegativity
• Ability to attract electrons • Trends in period: • Increase left to right • Metals have lower electronegativity than nonmetals • Trends in group: • Decrease top to bottom IV. Ionization Energy
• Amount of energy needed to remove the most loosely bound electron from an atom • Trends in period: • Increase left to right • Strong nuclear charge- More energy needed to remove electrons • Trends in group: • Decrease from top to bottom • Electrons farther away- easier to lose, less pull from nucleus V. Reactivity
• Noble gases are unreactive- full valence shells • Groups I and II are the most reactive metals • Never found pure in nature and lose electrons easily • Group 17- most reactive nonmetals (gain electrons ea sily and never found pure in nature)
Smith 22 VI. Metallic vs. Nonmetallic Characteristics
• Metallic increases down a group, decreases down a period • Nonmetallic decreases down a group and increases across a period • Thin black line spirits metals and nonmetals • Metalloids (semimetals) have properties of metals and nonmetals Solubility
• Solvent- Substance that is present in the larger amount which does the dissolving (the solvent dissolves) • Water is the universal solvent (dissolves most things) • Solute- Substance that is present in smaller amount which gets dissolved I. Types
• Solid solutions- mixture of 2 or more solids (Sand, metal alloy, Bronze (Zn+Cu) • Gas solutions- mixture of 2 or more gases (Oxygen, nitrogen, hydrogen, carbon dioxide) • Liquid solution • Solid in water (Salt water) • Gas in liquid (Carbonated dranks) • Liquid in liquid (water and juice) • Miscible- Liquids that mix in any amount (water and wine) • Immiscible- Liquids that cannot mix in any amounts (oil and water) • Aqueous • Electrolyte- Salt or ionic compounds that when dissolved in water will conduct electricity • Non-Electrolyte- Compound when dissolved in water that won’t conduct electricity (Any covalent bond, like sugars) II. Properties
• Homogenous mixtures • Clear, do not disperse light • May have color (transition elements) • Solute will not settle out
Smith 23 • Will pass through a filter (filtration cannot separate) III. Factors
• Nonpolar Molecules (fats) • Solvent- Nonpolar • Nonpolar molecule is soluble • Molecules will mix together • Polar solvent (Water) • Nonpolar molecule is insoluble • No attraction between the molecules • Polar Solute • Like Alcohol • Nonpolar solvent (carbon tetrachloride) is insoluble (no attraction between molecules) • Polar Solvent (Water) is soluble (attraction between molecules) • Ionic Solute • Solvent is nonpolar: insoluble and ions formed cannot attract to anything in the solvent • Solvent is polar: soluble and ons attract to the positive and negative ends of the solvent (H2O) • Temperature • As temperature increases: • Solids become more soluble in water and is true in most cases • Gases’ solubility decrease in liquid (think of soda) • Pressure • As pressure increases, gases become more soluble and has no effect on solids • To make gases soluble: High pressure and low temperature
Smith 24
Formulas and Equations Chemical Equations
• A chemical equation represents the starting and ending materials as: • Reactants !
Products
• The arrow represents yields or produces • The reactants are the starting material • Products are the ending material • Example: C+O2 ! CO2 Or Carbon + Oxygen yields Carbon Dioxide • Carbon and Oxygen are the reactants, Carbon Dioxide is the product Endothermic Reaction
• Energy is added for reaction to occur and is a reactant • Example: AB+Energy ! A + B • Example: Photosynthesis (Sunlight is added) • 6CO + 6H O + Energy 2 2
!
C6H12O6 + 6O2
• Since energy is absorbed in an endothermic reaction, the surrounding environment will decrease in temperature Exothermic Reactions
• Energy is removed/released for a reaction to occur and is a product • Example: A+B Yields to AB + Energy • Example: Cellular respiration (energy is released as heat, ATP) • C H O + 6O 6 12 6 2
!
6CO2 + 6H2O + Energy
• Since energy is released in an exothermic reaction, the surrounding environment will increase in temperature Law of Conservation of Mass
Smith 25 • Matter (mass) is neither created nor destroyed • In a chemical equation, both sides of the arrow must have the same amount of each type of atom • Example: H2 + O2 Yields to H2O • This equation is not balanced since there are 2 H and 2 O on the left but only 2 H and 1 O on the right • To balance: Coefficients will be added • Must be whole #s • They will apply to all elements within the formula • And can only be but before the entire formula, not in between • Subscripts may never be changed • All diatomic (HOFBrINCL) elements must be written as diatomic when they are not combined with any other element • If polyatomic ion is present on both sides of the equation, you may treat it as one thing or separate out the elements within it. • Example: Ca(NO3)2 : you have either 2 nitrates or 2 nitrogen atoms and 6 oxygen atoms Types of Chemical Reactions
• Synthesis: 2 or more reactants making only 1 product • A + B ! AB • 2H2 + O2 ! 2H2O • This type of bond is always exothermic due to the bond formation • Decomposition/Analysis Reaction: 1 Reactant making 2 or more products • Opposite of synthesis • AB ! A + B • 2H2O ! 2H2 + O2 • This type of reaction is always endothermic due to the breaking of bonds
Smith 26
• Single Replacement/Displacement: When a single element switches places with anothe r element in a compound • Always 2 reactants (one single element + one compound) and always 2 products (one single element + one compound) • A + BC ! B + AC • Mg + 2HCl ! H2 + MgCl2 • Double Replacement/Displacement: When 2 elements switch places with other e lements • Always 2 reactants (2 compounds) and always 2 products (2 compounds) • Outer 2 pieces come together and inner 2 pieces come together • NaOH + HCl ! NaCl + HOH (H2O) Identifying Reactions
• Single Replacement • Not all reactants will react • To determine if a single replacement reaction will occur: • Determine if the single element on the reactant side is a metal or nonmental • Then find this element on Table J and compare it to the corresponding metal or nonmetal of the other reactant • If the single element is higher on Table J than the metallic/nonmetallic element in the compound then the reaction will occur • Being higher means that the element is more reactive and can therefore “replace” the other • Double Displacement • 3 Situations in which a double replacement reaction will occur between two aqueous ionic compounds • 1: If one of the products is insoluble (It doe sn’t dissolve and therefore forms a precipitate, and the other is soluble • AgNO3 (aq) + NaCl (aq)
!
AgCl (s) + NaNO3 (aq)
Smith 27 • 2: If one of the products is a gas and the other product is aqueous • Na2S(aq) + 2HCl(aq) ! H2S(g) + 2NaCl(aq) • 3: If one of the products is water and the other is aqueous • NaOH (aq) + HCl (aq) ! H2O (l) + NaCl (aq) • A double replacement reaction will not occur if both products are a queous Unknown Reactants and Products
• You may have to predict the formula of an unknown reactant or product • Example : 2Na + 2H2O ! X + 2NaOH • To find X: • Tally the amount of atoms you have on either side of the arrow • 2 Na 2 • 4H2 • 2O2 • It seems that we are missing 2 H atoms • If written correctly, we can find the missing piece (H2) • 2Na + 2H2O !
H2 + 2NaOH
• If it’s a compound that is missing: • Write the symbols of the element • If ionic: Criss-cross charges to get exact formula • If covalent: think of what common molecule it could be Kinetics
• Studies the rates of chemical reactions and how quickly they occur Collision Theory
• For a reaction to occur, reactant particles must collide with • Enough energy (activation energy) • Correct spatial orientation • When these conditions are met, an effective collision has occurred • One that makes new products
Smith 28 • In general, as the number of effective collisions between particles increases, so does reaction rate Factors Affecting Reaction Rates
I. Nature of the Reactants • Covalent or organic (containing C) substances react slower than ionic • Because they have more bonds that must be broken as they react • Ionic substances react more quickly • Have no true bonds • Ions held together by electrostatic force II. Concentration (Mol/L) • Means how much stuff per Liter of space • If concentration is increased, then more collisions between particles will result • More collisions=faster reaction rate • In general: increase concentration, increase reaction rate III. Surface Area • Exposing more of a reactant’s surface area will lead to faster reactions • This is because there will be more reactant particles contacting the other reactants • Surface area can be increased by breaking down a chunk of a reactant into smaller pieces or powder • A powdered form of a substance always gives the most surface area IV. Pressure • Only affects reactions involving gases • Increasing pressure increases the concentration of a ga s • More in less space or smaller volume • This results in a faster reaction rate • (Only affects gases of different moles) V. Temperature • Measure of the average kinetic energy of particles • Higher temperatures allow particles to move faster with more energy
Smith 29 • The faster they move, the greater the chance of them colliding • Increased collisions leads to faster reaction rate • In general, increased temperature=increased reaction rate VI. Catalysts • Substances that increase the rate of a reaction • By providing a quicker, alternate pathway that requires lower activation e nergy • These substances are not changed in any way throughout the reaction • Examples: enzymes, heavy metals (Pt) Role of Energy in the Reaction
I. Activation Energy • Needed to start a reaction • Varies based on the nature of the reactants • Energy can be absorbed (endothermic) or released (exothermic) in a chemical reaction II. Enthalpy • Heat of the reaction (!H) • This is the difference between the energy of the produ cts and reactants • !H+HP-HR • Free Element- H=0 (J) Free element has 0 heat • Heat of Formation- Amount of heat released or consumed when 1 mole of a compound is produced from the free elements • !H= !HProducts -H!Reactants • Driving Forces- Low Enthalpy (- !H), High Entropy (+ !S)
• Free Energy- Combined Effect • !G= !H-(T)(!S) • -!G=Decrease in free energy of the system + a favorable drive • + !G= Increase in free energy of the system + a n unfavorable drive • Exothermic Reactions • Energy is released • Energy is a product
Smith 30 • Products have less energy than reactants • !H is negative (This means that energy is released by the reactants, not a negative amount of energy) • Table I • Endothermic Reactions • Energy is absorbed • Energy is a reactant • Products have more energy than reactants • !H is positive (this means that the reactants have ab sorbed energy)
Potential Energy Diagrams •
• Activated Complex- Transition state where reactants either become products or reform reactants • A catalyst increases the reaction rate by lowering activation energy
Smith 31 • Causes the activated complex and activation energies to be lower, but does not change the head of the reaction • If the potential energy diagram is endothermic, the p roducts will be higher than the reactants and have more energy • If the potential energy diagram is exothermic, the produc ts will be lower than the reactants and will have less energy Entropy !S
• Measure of randomness or disorder • As entropy increases, !S becomes more + • As entropy decreases, !S becomes more • Physical Changes • Phase Changes- Endothermic processes and when a substance dissolves, entropy increases • Free elements are less stable and have more entropy than compounds • Increase T increases entropy • When 2 different gases mix, entropy increases Spontaneous Reactions
• There is a tendency in nature to favor • Exothermic Reactions • More stable with less energy • Higher entropy • Easier to be disorderly than orderly Equilibrium
• Most reactions can occur in both the forward and reverse directions • Both reactions will occur at the same rate • This means that the forward reaction (making products) is the same as the reverse (reforming reactants) • It Does NOT mean that the concentration (amounts) of the same reactants and products are equal • The concentration of the reactants and products are constant
Smith 32 • Equilibrium will only occur if nothing leaves the system • Example: If gas escapes or solid is formed (No eq uilibrium) • System must be closed Phase Equilibrium • Occurs when a substance is changing its phase of matter • At the melting point • For a short period of time, both the solid and liquid phases of matter are in equilibrium with each other • At the boiling point • For a short period of time, both the liquid and gas phases are in equilibrium Solution Equilibrium • In a saturated solution, the rate of dissolving equals the rate of recrystallization • A closed soda bottle is also at equilibrium • The rate of the CO2 dissolving in the soda equals the rate of the dissolved CO2 escaping • If pressure is increased on the system, the reaction will shift left and more of the CO2 will stay dissolved • If pressure is decreased on the system, the reaction will shift right and more of the CO2 will escape as gas Le Chatelier’s Principle • A system at equilibrium can be disturbed by placing a stress on it • Include: Change in temperature, concentration, pressure • When a system has a stress placed on it, the reaction shifts to relieve the stress and reestablishes equilibrium • Concentration Changes: • When the concentration of either a reactant or product is • Increased • Reaction shifts Away from the substance increased • Decreased • Reaction shifts toward the substance decreased
Smith 33 • Temperature Changes • If you increase temperature, the endothermic reaction is favored and shifts away from heat • If you decrease temperature, the exothermic reaction is favored and shifts toward heat • Pressure Changes • Only affects gaseous substances • When pressure is increased, reaction shifts toward less moles • When pressure is decreased, reaction shifts toward more moles • When # of moles is the same on both sides, the pressure has no effect • To determine the # of moles, add the coefficients on the left side and the right side and compare • Whatever you do, the reaction does the opposite Oxidation-Reduction/Redox Reactions
• A redox reaction is: • A type of chemical reaction • Both reduction and oxidation occur simultaneously due to a competition for electrons between atoms • Reduction: Gain of Electrons (GER) • A species (element or ion) gains electrons • Plating (metal spoon plated with silver) • Oxidation: Loss of Electrons (LEO) • A species loses electrons • Corrosion (car rusting - iron is losing electrons) • LEO says GER Oxidation Numbers
• Found in the upper right hand corner of each element on the reference table (similar to charge) • Describe the number of electrons gained or lost by an atom • Rules for assigning numbers • In an uncombined element, Ox # is 0 • Certain metals only have oxidation #
Smith 34 • Group 1= +1 • Group 2= +2 • Fluorine is always -1 in compounds • The other halogens (Cl, I, Br) are also -1, but only when they are the most electronegative element in the compound • Hydrogen is +1 in compounds unless it is combined with a metal • If it’s with a metal, it is -1 • Oxygen is usually -2 in compounds • There are exceptions • The sum of oxidation numbers in all compounds must equal zero • The sum of the oxidation numbers in polyatomic ions must equal the charge on the ion Half Reactions
• Oxidation and Reduction occur simultaneously • One cannot occur without the other • During redox, there is always a conservation of mass and charge • There are 2 half-reactions that occur in redo • Reduction - Where electrons gained are placed on the left hand side of the arrow (reactant) • Oxidation - Where electrons are placed on the right hand side of the arrow (product) Electrochemical Cells
• In redox reactions, there is a chemical reaction and an exchange of electrons between the particles being oxidized and reduced • Electrochemical Cell: Involves a chemical reaction and a flow of electrons • Voltaic: Named after Alessandro Volta, and is an electrochemical cell in which a spontaneous chemical reaction produces a flow of electrons • Electrolytic: Requires an electric current to force a non -spontaneous chemical reaction to occur • Have two surfaces called electrodes (An Ox, Red Cat) • Electrode: Site at which redox occurs • Anode: Electrode where oxidation occurs
Smith 35 • Cathode: Electrode where reduction occurs
Spontaneous Reactions- Voltaic Cells
• If a strip of zinc is placed into a solution of lead nitrate, the zinc will be oxidized and the copper ions will be reduced • Zn(s) + Cu2+(aq) -> Cu(s) + Zn 2+(aq) • Zn(s)->Zn2+(aq) + 2e• Cu2+(aq) + 2e- -> Cu(s) • In a voltaic cell, a salt bridge connects the two containers and provides a path for a flow of ions between the beakers • In such voltaic cell, when a strip of zinc is located in one beaker and copper ions are in solution in another beaker, the reaction can occur as if the solution were in the same beaker • When electrons are lost during oxidation at the anode, the travel through the wire to the cathode • The material being reduced gains electrons • E0 Cell= E0Reduction-E0Oxidation Non-Spontaneous Reactions- Electrolysis
• Electricity is used to force a chemical reaction • Used to obtain active elements such as sodium and chlorine by the electrolysis of fused (molten cells) • 2NaCl (l) -> 2Na(s) + Cl 2(aq) • Used to electroplate metals onto a surface • Have several things in common with a voltaic cell: • Both use redox reactions • The anode is the side of oxidation • The cathode is the site of reduction • The electrons flow through the wire from anode to cathode Acids, Bases and Salts I. Theories
I. Arrhenius (1887) • Acid - H+ is released (H 2SO4, HCl) • Bases- OH- is released (NaOH, Ca(OH) 2 II. Br ønsted-Lowry (1923) • An acid donates a H+ (proton) and is a proton donor • A base accepts a H+ (proton) and is a proton acceptor • HF + H2O -> F- + H3O+ ---HF is acid, F is conjugate base, H 2O is base, H 3O is conjugate acid.
Smith 36 • Conjugate acid is what is formed after a base gains a H+ ion • Conjugate base is what remains after the acid donates its H+ ion. • Base + Acid -> Conjugate Acid + Conjugate Base • Acid becomes conjugate base • Base becomes conjugate acid III. Lewis Theory (1923) • An acid is any substance that accepts a pair of electrons (2e -) • A base is any substance that donates a pair of electrons • BF3 + :NH3 -> H3 N:BF3 : = 2 free electrons • BF3 - Lewis Acid • :NH3 - Lewis Base • H3 N:BF3 - Product • Lewis - contains Bronsted and Arrhenius • Bronsted- Contains Arrhenius • Arrhenius- Does not contain Bronsted and Lewis • A substance that is an acid or base under Arrhenius theory is also an acid and base under BronstedLowry Theory-Each succeeding theory is more inclusive. • Aq solutions of acids conduct electric currents (electrolytes) (strong acid=good conductor) • Strong base=good conductor and a weak base=bad conductor • Polar covalent aids when dissolved can conduct electricity II. Amphoteric Molecules
• Amphoteric molecules are molecules that act as an acid or base depending on what it is mixed with • HCl + H2O -> H3O+ + Cl- Water acts as the base • Water is composed of both H+ and OH• NH3 + H2O -> NH4+ + OH- Water acts as the acid • Water can ionize H 2O (acid) + H2O (base) -> OH- + H 3O+ • [H3O]+ = [OH]- Concentration=Concentration • K eq= Concentration of Products/ Concentration of Reactants
[H 2O]= 55.6 Mol/L
• K eq = [H3O+][OH-] / [H2O][H2O] • K eq= [10-7][10-7] / [H2O]2 • K eq= (55.6)2 = [H3O+][OH-]= [10-7][10-7]= 10-14 (Where pH comes from) • K W= K eq(55.6)2 = Constant of water • K w=[H+][OH-]= 10-14 III. Ph Scale=Power of Hydrogen
• 0-14 (Measures H+ Concentration) • pH = -log(H+)
Smith 37 • 7- Achieved in neutralization reaction of H +OH- (Neutralization results in production of water and salt) • H+ + OH- -> H2O Neutralization reaction • High pH- Greater # of OH - ions -> pH+pOH=14 • Lower pH- Greater # of H + ions • If -log(H+)=5, what is the pH of the base? • -log(1x10 -5) = pH of 5 14-5= 9 pOH- 9 • Scale based on power of 10 • Ph of 1 is 10x more acidic than 3 • 14 is 100x more basic than 12 • If pH changes from 6 to 2, what happens? 6-2=4 10 4= 10,000 more acidic IV. Titrations
• The process of adding measured volumes of an acid or base of known concentrations to an acid or base of unknown concentration until neutralization occurs • Performed to determine the concentration of unknown solution • The solution of known concentration is called the standard solution V. Titration Equation
• MAVA=MBVB • MA= Molarity of acid/ H+ • VA= Volume of acid • MB= Molarity of base/OH• VB= Volume of base • In titration (neutralization), must be 1:1 ratio between H+ and OH-
Nuclear Chemistry
I. Stability of Nuclei • Ratio of protons and neutrons that determines stability • Atomic numbers greater than 83 are radioactive (Unstable isotope=radioisotope) • When an unstable nucleus decays, it emits radiation in the form of alpha/beta particles, positrons/gamma radiation. • K-Capture Process- When nucleus captures an electron from 1st energy level, nucleus will be unstable: there is spontaneous decay. • Alpha Particle- Helium nucleus, 2p, 2n, (+), low penetrating power • Beta Particle- Electron whose source is an atomic nucleus (-), moderate penetrating power • Positron- Identical to electron, but with positive charge
Smith 38 • Gamma- Similar to X rays, but greater energy - not deflected by electric field, high penetrating power II. Alpha Decay • Unstable nucleus emits alpha particle- nucleus is alpha emitter • Characteristic of heavy nuclei • As nucleus emits alpha particle, atomic # decreases by 2, and mass # decreases by 4 III. Beta Decay • Nuclear disintegration from electron- undergoes beta decay and is a beta emitter • Emission of electron during conversion of neutron to proton (1/0 n -> 1/1 p + 0/-1 e) • When a nucleus emits an electron, the charge of the nucleus increases by 1, atomic number increases by 1 IV. Positron Emission • Production of positron during conversion of proton to a neutron • When a nucleus emits a positron, the charge of the nucleus decreases by 1, thus the atomic number decreases by 1 V. Nuclear Equations • Mass and charge MUST balance on both sides (14/7N + 4/2He -> 17/8O + 1/1H) • # of charge of reactants= 9 • # of mass#=18 • Concept of conservation of charge and mass number is used to identify particle VI. Transmutation • When the nucleus of one element is changed into the nucleus of another • Can be either natural or artificial • Natural: One reactant only • Artificial: Two reactants and occurs by bombarding the nucleus with high energy particles or by colliding a nucleus with a neutron • Fission: Reaction that splits a heavy nucleus to produce lighter ones (Captures a neutron and becomes unstable)
Smith 39 • Fusion: Occurs on sun, combines with light nuclei to form heavier ones (Hydrogen nuclei react in a series to produce helium nuclei). Does not occur on Earth because of the extremely high temperatures and pressures needed
Organic Chemistry
I. Bonding of Carbon Atoms • The ability of C to form many different compounds is based on the tendency to covalently bond with other C atoms • One single bone = Saturated • Sharing two e- : double covalent bond - Unsaturated • Aliphatic Hydrocarbons: hard carbon atoms linked in chain • Aromatic Hydrocarbon: Contains one or more benzene rings
II. Aliphatic Hydrocarbons • Alkanes: Single covalent bond- Saturated C NH2N+2 • Alkenes: Double covalent bond- Unsaturated C NH2N • Alkynes: Triple covalent bond- Unsaturated C NH2N-2
II. Hydrocarbon Radical • A hydrocarbon molecule from which a hydrogen atom ha s been removed
->
---> One less hydrogen= Radical (MethyL)
III. Hydrocarbons • Homologous- group of related compounds in which each member differs from the one before it by the same unit
Smith 40 • Alkane- Release energy when burned (CH4, C2H6, C4H10) (as # of C increases, so does the boiling point because of the amount of bonds) • Alkenes- 1 Double bond- provide chemists to make other materials- most important is ethane, ethylene (forms plastic) • Alkyne- Unsaturated hydrocarbon that contains triple bond (Ethyne, acetylene, fuels welding torches) IV. Isomers • Same molecular formula, but different structural formula • MethyL propane C4H10 • Butane- has 2 isomers • 2,2 Di-MethyL Propane• (3 Radicals)
V. Alkane- Paraffins • C NH2N • Methane, Ethane, Propane, Butane, Pentane • Ends in ANE • Single Bond
VI. Alkenes- Olefins- Ethylene Series • C NH2N • Ends in ENE • Ethene, Propene, Butene, Pentene • Double Bond
VII. Alkynes- Acetylene Series • C NH2N-2 • Ends in YNE • Ethyne, Propyne, Butyne, Pentyne
Smith 41 • Triple Bond VII. Organic Reactions • Substitution: Reactions in which a H atom of a hydrocarbon is replaced by another atom or group (Exists only between alkanes)
• Additions: Reactions in which one bond of a double bond is broken so that atoms may be added to the hydrocarbon (Will also occur with one or two bonds breaking in a triple bond) (second class alkenes)
• Elimination- Reactions in which atoms are removed from A hydrocarbon to create a double or triple bond
• Esterification: formation of an ester by reacting an alcohol With an organic acid and removing H2O
Smith 42
• Hydrogenation (additions): The addition of H atoms when a double or triple bond is broken
• Combustion: Alkanes burn in air to produce carbon dioxide and water vapor
• Cracking: Process by which complex organic molecules are broken into simpler molecules; involves heat or heat and a catalyst C3H8 (460C ->) C2H4 + CH4 • Polymerization: Many single units (called monomers) join together to make a polymer (breaking double and triple bonds)
