Chemistry Review

Prefixes Tera Giga Mega Kilo Hecto Deca Deci Centi Milli Micro Nano Pico

Problem 1.1: 3.45×1016 + 9.80×1018 __________________

Meanings Trillion Billion Million Thousand Hundred Ten Tenth Hundredth Thousandth Millionth Billionth Trillionth

make exponents EQUAL first

18

Exponents 1012 109 106 103 102 101 10-1 10-2 10-3 10-6 10-9 10-12

Problem 1.2: 6.12 × 10-2 - 4.9 × 10-4 ___________________

.0345×10 + 9.80 ×1018

6.120 × 10-2 - .049 × 10-2

__________________

___________________

9.8345×1018

Measurement Weight Time Length Mass Temperature Energy Speed Area Volume Density Molarity

SI Unit Newton Second Meter Kilogram Kelvin Joule

Derived Units

m/s m2 m3 kg/m3 moles/L

Significant Figures: any number not zero zero between two other numbers zero as the last digit and RIGHT of the decimal zero as a placeholder

6.071 × 10-2

Mole

Number City atomsmoleculesFormula units

Families tell you the number of valence electrons. Periods tell you the number of shells.

Gram City

All atoms of every single element are not necessarily stable, but they are electrically neutral. To name a compound from its chemical formula: Name the first element, and then the second, but change the ending to –ide. Ex. NaBr = Sodium Bromide When using a conversion factor, you always use ONE mole.

Steps to Calculate Molarity: 1. convert units to moles if needed 2. put in the form of moles per liter 3. done

Preparing a Certain Molarity: How Much of What Kind and you want to get GRAMS

Oxygen one atom – 16 AMUs one mole – 16 grams Pg. 205 #42: 5.23 g Fe(NO3)2 in 100.0 cm3 of solution 5.23 g  moles: 1 mol 5.23 g × ----------- = .029080 mol 179.845g

____________________

179.845

.029080 mol .029080 mol ----------------- = ---------------- = .029080 M 100 ml 1L Pg. 205 #46: 1.00 dm3 of 3.00M NiCl2 How Much of What Kind and GET GRAMS 3 mol 129.599 g 1.00 L × --------× ------------- = 388.797 g NiCl2 1L 1 mol

Percentage Composition: pg. 208 #51: aluminum sulfide, Al2S3, percentage composition Al2 = 53.964 53.964 S3 = 96.3 ------------ = 35.91% Al Al2S3 = 150.264 150.264 Steps 1. convert grams to moles 2. write formula using moles 3. divide by smaller number and round to nearest whole

1(Fe) - 55.845 2 (N) - 28.000 6 (O) - 96.000

96.3 ----------- = 64.08% S 150.264

Empirical Formula pg. 210 #56: empirical formula of a compound that contains 1.67 g Ce, and 4.54 g I.

1.67 x 1 mol/140.12 g = .011918 4.54 x 1 mol/126.90 g = .035776 Ce.011918I.035776 = CeI3

%Ce = 26.90% %I = 73.1%

Hydrates Steps 1. convert grams to moles 2. divide by smaller number and round to nearest whole 3. write equation *substance is ALWAYS ONE.

BaI2 · ? H2O; mass of 10.407 g; dry BaI2 has a mass of 9.520 g. 9.520 BaI2 · .887 H2O

9.520 BaI2 x 1 mol/391 g = .0243mol / .0243  1 .887 H2O x 1 mol/18 g = .0493mol / .0243  2

BaI2 · 2H2O

Reaction Synthesis Decomposition Single Displacement Double Displacement Combustion

Description Substance + Substance  Compound Compound  Substance + Substance Element + Compound  Different Element + Different Compound Compound + Compound  Compound + Compound Hydrocarbon + Oxygen  Carbon Dioxide + Water 1. Carbon 2. Hydrogen 3. Fix Oxygen (Half Molecules May Be Needed)

Diatomic Molecules: H, N, O, F, Cl, Br, I (make a 7 on Periodic Table) Nitric Acid – H(NO3) Carbonic Acid – H2(CO3) Sulfuric Acid – H2(SO4) Multiple Oxidation States: Copper, Lead, Iron, Nitrogen (l) – liquid (g) – gas (c) (cr) – solid (aq) – aqueous (in water) anode/cation – positive (anode – ain’t negative) cathode/anion – negative

Q = m(ΔT)Cp Pg. 72 #35: BP: 2600°C & FP: 1400°C, Cp=0.1733 J/g°C Joules to heat 10.0g from FP to BP? Q = 10.0g (1200°C).1733 J/g°C Q = 2079.6 J OR 2.0796 kJ Ex: 10g Ag ring @ 50°C put into 100 ml of H2O @ 1°C; find final temperature of the system QLOST BY SILVER = QGAINED BY WATER m(ΔT)Cp = m(ΔT)Cp hot final cold 10(ΔT).2350 = 100(ΔT)4.18 [50°C - (X] - 1°C) 10(50 – x).2350 = 100(x – 1)4.18 2.350(50 - x) = 418(x - 1) (50 – x) = 177.87(x – 1) 50 – x = 177.87x – 177.87 227.87 = 178.87x x = 1.27°C Distillation – Boiling Points

Crystallization – Solubilities

Mass to Mass: pg. 237 #28: C6H12O6 + 6O2  6CO2 + 6H2O calculate oxygen, 12.5g glucose 6(12) – 72 \ 180 192 AMU 12(1) – 12 = 180 ------ = -------------X = 13.3 g O2 6(16) – 96 / 12.5 X Mass to Heat Steps: 1. find H°f prod and H°f reac USE COEFFICIENTS!!! 2. plug into ΔH°f rxn = H°f prod – H°f reac 3. use Mass to Heat proportion

pg. 241 #35: NH3 + HBr  NH4Br + (–) 188.32 kJ compute heat of reaction for 193g NH4Br -46.11 + -36.40  -270.83 NH4BR AMU = 98 ΔH°f rxn = H°f prod – H°f reac 98 193_ -36345.76 = 98x = = -188.32 kJ -188.32 X x = -370.875 kJ Mass to Mass: Mass to Heat: AMU known AMU unknown ΔH°f rxn = H°f prod – H°f reac ------------------- = -------------------AMU g Known g grams known X ------------- = ------------X = grams of unknown ΔH°f rxn X (heat) Democritus – “atomos” means indivisible Lavoisier (La-voice-ee-ay) – Law of Conservation of Mass Joseph Proust – Law of Definite Proportions Specific substances will always contain elements that appear in the same ratio by mass. (subscripts) John Dalton – Law of Multiple Proportions Combining masses of one element with a constant amount of another will do so in a RATIO of small whole numbers. (coefficients) Know the difference between LIMITING and EXCESS substances. Gay-Lussac – Law of Multiple Proportions in respect to Gases. At constant temperature and pressure, volumes of reacting gases and their products will be in a ratio of small whole numbers. Amadeo Avogadro – Under identical conditions of temperature and pressure, equal volumes of gases contain equal numbers of particles. (Particles are so far apart that particle size becomes irrelevant.) At S.T.P. – one mole of any gas has a volume of 22.4 L. (to work a problem with this, convert grams  moles, use a mole to mole ratio, and convert moles  L) Dalton’s Hypothesis – formed the basis for our modern atomic theory 1. All matter is made up of atoms. 2. Atoms of the same element are alike. INCORRECT (He didn’t envision ISOTOPES) 3. Atoms of different elements are different. 4. Atoms of different elements can combine in simple, small, whole number ratios to form compounds. anode/cation – positive (anode – ain’t negative) cathode/anion – negative Crooks Tube – cathode rays seen streaming from the cathode (later found to be electrons) J.J. Thomson – discovered the electron (NOT the charge, just the RELATION TO MASS) – discovered the proton – 1836 times more massive than the electron Millikan and the “Oil Drop Experiment” – proved that electron charge is negative by using an atomizer. Neutron

– Walter Bothe – predicted discovery – Rutherford – contributed evidence – Chadwick – credited with discovery

Atomic Number – Z Mass Number – A A – Z = number of neutrons In a stable element, the number of electrons and protons will be equal. Moseley – determined several elements’ atomic number by using X – rays. Subatomic Particle Mass Charge Proton 1 AMU 1+ Neutron 1 AMU Neutral Electron 0 AMU 1-

Mass Spectrometer – separates different isotopes of a single element Rutherford Gold Foil Experiment proved: 1. Nucleus is a small, dense structure at the center of an atom. 2. Nucleus has a positive charge. 3. Nucleus is surrounded by an electron cloud that is very distant from the nucleus. Neils Bohr – planetary model of the atom (definite energy levels, but not on the same plane) Electromagnetic Spectrum – Wavelength (λ); Frequency – cycles/second – Hz; Amplitude - height Max Planck – quantum theory (energy is given off in little “packets” called “quanta” or photons.) Newtonian Mechanics vs. Quantum Mechanics ordinary objects at ordinary velocities vs. infinitesimally small objects at light speeds De Broglie (de BROY) – Wave-Particle Duality of Nature Refraction – bends due to change in medium; the more bends, the easier something is to see Index of Refraction – degree to which a substance bends light Momentum = mass x velocity Werner Heisenberg – Heisenberg Uncertainty Principle – We can NEVER know both the exact position and momentum of an electron. The more certain we are of one, the less certain we are of the other. – WHY: To determine the position of an electron, you have to see it. If you see an electron, it has had a collision with a photon of light. The collision alters the electron’s velocity, thus changing its momentum. Schrödinger – developed wave equation (describes electron behavior) Quantum Numbers 1. Principal Quantum Number – N – cloud size and energy level To determine the maximum number of electrons in an energy level = 2n2 (n - # of energy levels) 2. Azimuthal Quantum Number – L – cloud shape and sublevels To determine the number of sublevels that make up an energy level: l = n (n - # of energy levels) 3. Magnetic Quantum Number – M – orbitals and direction in space of orbitals Orbitals make up sublevels and can house up to two electrons. Sublevel Number of Orbitals Max eS 1 2 P 3 6 D 5 10 F 7 14 odd numbers x 2 4. Spin Quantum Number – S – spin of electron (one spins clockwise, one spins counter-clockwise) s p d f 1 2 3 4 Half Full, One CLOCKWISE + 1 --2

Full, One CLOCKWISE, One COUNTER-CLOCKWISE + 1 --__ 2

Electron Configuration – Diagonal Rule Chart: (Go through one, one, then two, two, three, three, four, four) 1s2 2s2 2p6 3s2 3p6 3d 10 4s2 4p6 4d 10 4f 14 (energy level) (sublevel) # of electrons 5s2 5p6 5d 10 5f 14 6s2 6p6 6d 10 7s2 7p6 (8s2) Ex: Bromine (35 e-) - 1s2 2s2 2p6 3s2 3p6 3d 10 4s2 4p5 Lewis Electron Dot Diagram – show the valence electrons as dots around the chemical symbol of the element. 6 3

4 7

Ga

2 1

5

8

Z = 31:

Ga

Pauli Exclusion Principle – no two electrons in the same atom can have all the same quantum numbers. Kinetic Theory explains the effects of heat and pressure on matter. Three Assumptions:  All matter is composed of small particles.  These particles are in a constant state of motion (even solids).  Collisions between particles are perfectly elastic (which means energy transfers from particle to particle but does not leave the system). Mean Free Path – distance a particle travels between collisions Pressure – FORCE/AREA Two Most Important Causes of Gas Pressure – Force and Frequency of Collisions

1 kPa = 7.5 mmHg Measuring Gas Pressure with a Manometer – open relies on atmospheric pressure to work – closed works independently of atmospheric pressure; referred to as absolute pressures Barometer is a closed manometer used to measure atmospheric pressure First manometer was invented by Torchelli. One kiloPascal = 7.5 mm

56 mm 7.5 mm ----------- = ---------7.5 kPa 1 kPa

CONSTANT

7.5 kPa + 102.2 kPa (atmospheric pressure) 109.7 kPa

The gas moved the mercury towards the open end, so the gas pressure is greater than atmospheric pressure. Hence, you add the atmospheric pressure in the end. In a closed manometer, the number you get in the proportion is your answer because atmospheric pressure is not involved.

KE = ½m(v2) – mass x velocity squared Two Most Important Factors of Velocity: Temperature & Mass (more mass, slower movement) Heat Transfer: • Solid – definite shape, definite volume; particles are in fixed positions and vibrate around fixed points. • Liquid – no definite shape, definite volume; particles slip and slide past each other and vibrate around moving points. • Gas – no definite shape, no definite volume; straight-line paths between collisions, totally random. • Plasma – superheated gas, electrons get stripped away from nuclei; like atom soup. Thermal Expansion – atoms and molecules expand when heated. Magnetohydrodynamics – study of plasmas. Boltzman Distribution Law – Most particles in a given system have a kinetic energy at or about the average, but some are above and some are below. Saturated – When a substance is in equilibrium with its vapor, the gas phase is saturated. Dynamic Equilibrium – Rates of change from liquid to gas and vice versa are equal. Vapor – Prefers to be a liquid at room temperature Gas – Prefers to be a gas at room temperature

Le Chatelier’s (SHOT-LEE-AY) Principle – If you take a system at equilibrium and place a stress on that system, it will respond so as to relieve the stress. (Stress can be a change in temperature, pressure, or concentration) Vapor Pressure – the tendency of a substance to move towards the gaseous state. Volatile Non-Volatile High Vapor Pressure Low Vapor Pressure Weak Intermolecular Forces Strong Intermolecular Forces Rapid Evaporation Slow Evaporation Low Critical Temperature High Critical Temperature Low Boiling Point High Boiling Point Ex: Alcohol, Dry Ice, Gasoline Ex: Glycerol (Lotion), Ethylene Glycol Hydrogen Bonding – Hydrogen atoms are not affected by Van der Waals forces, but still have strong bonds. Melting Point – (same as freezing point), temperature at which a solid changes to the liquid state and the V.P. of the solid is equal to the V.P. of the liquid. Sublimation – when a solid changes directly to the gaseous state without passing through the liquid state. (very high vapor pressure) Ex: solid CO2 (dry ice), iodine, mothballs, snow Deposition – the process of changing directly from a gas to a solid without passing through the liquid state. Boiling vs. Evaporation – Evaporation occurs only at the surface and will happen at temperatures lower than the boiling point. Boiling Point – the temperature at which the vapor pressure on the surface of the liquid equals the vapor pressure within the liquid. Normal Boiling Point – the temperature at which a substance boils when the pressure on the surface in 1atm (101.325 kPa) Pressure will change the boiling point and the melting/freezing point. Liquefaction – the condensation of a substance that is typically a gas. To liquefy a gas, lower the temperature and raise the pressure. Tc (Critical Temperature) – the temperature above which no amount of pressure can cause liquefaction. (Tc is usually very high.) Pc (Critical Pressure) – the quantity of pressure required to liquefy a gas that is at its critical temperature. Phase Diagram – shows the relationship between temperature, pressure, and the physical state of a substance.

**Any point on any one of the three lines represents an equilibrium between two states of matter.**

Hf (Enthalpy of Fusion) – the quantity of heat that is required to melt one gram of a substance that is at its melting point. Hv (Enthalpy of Vaporization) – the quantity of heat that is required to vaporize one gram of a substance.

Ex: How much heat is needed to change the temperature of 10g of ice from -10°C to 110°C? Cpice = 2.06 J/g°C Hfice = 334 J/g CpH2O = 4.18 J/g°C HvH2O = 2260 J/h Cpsteam = 2.02 J/g°C • -10°C to 0°C Q = m(ΔT)Cp = 10 (10) (2.06) = 206 J • Hmelt = m(Hf) Q = 10 (334) = 3340 J • 0°C to 100°C Q = 10 (100) (4.18) = 4180 J • Hboil = m(Hv) Q = 10 (2260) = 22600 J • 100°C to 110°C Q = 10 (10) (2.02) = 202 J Total = 30,528 J S.T.P. – Standard Temperature and Pressure - 0°C (273K) and 101.325 kPa An Ideal Gas is composed of Point Masses – infinitely small points with no volume, only mass, and not subject to Van der Walls forces.

Boyle’s Law – (pressure change) If the temperature remains constant, the pressure will vary inversely as the volume. Ex: A gas has a volume of 200 cm3 and has a pressure of 150 kPa. What pressure will this gas have at a volume of 225 cm3? (Assume temperature is constant.) P1 = 150 kPa P2 = ___ kPa 200 cm3 T1 = 150 kPa x ----------- = 133.3 kPa T2 = 225 cm3 3 V1 = 200 cm V2 = 225 cm3 V↑ P↓ Pressure is going DOWN, so put the SMALLER one on top. Charles’ Law – (temperature change) the volume of a gas, held at a fixed pressure, will vary directly as the KELVIN temperature. (Kelvin because Celsius scale uses a zero and zero can mess you up in fractions.) Problems work the same as Boyle’s Law.

°C + 273 = K Combined Gas Law – You can correct a volume of gas for not only a change in pressure but also a change in temperature at the same time. Ex: A gas has a volume of 250 cm3, a pressure of 103 kPa, and a temperature of 75°C. What volume will this gas have if the pressure is 250 kPa and the temp is 120°C? P1 = 103 kPa P↑ V↓ P2 = 250 kPa 103 kPa 393 K 3 T1 = 348 K T↑ V↑ 250 cm x ------------- x --------- = 116.32 cm3 T2 = 393 K 250 kPa 348 K 3 V1 = 250 cm V2 = ___ cm3

VA mB ---- = ---VB mA

KE = ½ m (v2) The relative rates of diffusions of two gases is inversely proportional to the square roots of their molecular masses. Always put the LESS MASSIVE molecule ON TOP on the LEFT.

Ex: Compare the rate of diffusion of Ammonia to that of Hydrochloric Acid. Ammonia – NH3 – 17 AMU Hydrochloric Acid – HCl – 36.5 AMU VNH3 mHCl 36.5 -------- = --------- = ------ = 1.5 NH3 travels 1.5 times faster than HCl VHCl mNH3 17 Ideal Gas Equation: PV = nRT (Pressure)(Volume) = (number of moles)(gas constant)(temperature in K) Ex: What is the pressure exerted by .622 mol of gas in a 9.22 dm3 vessel at 16°C? CONVERT TO DM3!!! 3 P(9.22 dm ) = .622 (8.31) (289) = 1493.79 P = 162 kPa Gas Collection by Water Displacement Dalton’s Law – The total pressure of a mixture is the sum of the partial pressures of the gases that compose the mixture. Partial Pressure – the pressure exerted by an individual gas. The partial pressure exerted by a gas is CONSTANT, whether it is the only gas present or in a mixture of gases. All gases in a mixture have the SAME VOLUME, therefore differences in their partial pressure is due to the numbers of molecules present. N2 PP = 75 kPa + O2 PP = 25 kPa Total PP = 100 kPa Dry Gas – the gas in the top of the container (H2 in this case) minus the water vapor. Ex: ↑ If the volume is 250 ml, the temperature is 22°C, and the vapor pressure is 90 kPa, what will the volume of the dry H2 gas be at standard pressure? V.P. = 90 kPa P1 = 87.4 kPa - V.P. of H2O = 2.6 kPa (at 22°C) P2 = 101.325 kPa V.P. of H2 = 87.4 kPa T1 = 87.4 kPa T2 = 250 ml x ----------------- = 215 ml V1 = 250 ml 101.325 kPa V2 = ___ ml P↑ V↓ When acids/bases combine with H2O, they conduct electricity and are called electrolytes. Arrhenius’ Theory – When these substances go into solution, they break down into charged particles. Gilbert Newton Lewis’ Theory – Not all acids and bases focus on proton transfer (Hydrogen). The Lewis Theory focuses on electron transfer and is very commonly associated with organic acids. Lewis Acid – electron pair ACCEPTOR. Lewis Base – electron pair DONOR. Acids – sour; lemons cation is usually Hydrogen anion either IS or CONTAINS a non-metal any substance which, when dissolved in water, will produce hydrogen ions. (H+ is a bare proton.) HCl  H+ + Cl- (ionization) Brönsted-Lowry Theory - any substance that will DONATE a proton. Bases – bitter; soap

cation is a metal

anion is usually Hydroxide (OH)1-

any substance which, when dissolved in water, will produce hydroxide ions. Na(OH)  Na+ + (OH)Brönsted-Lowry Theory - any substance that will ACCEPT a proton. Salts – salty; table salt Conjugate Base – particle that remains after the acid has donated a proton. Conjugate Acid – particle formed after the base gains a proton from the acid. HCl + H2O  H3O+ + ClNH3 + H2O  NH4+ + OHBinary Acids – contain two and only two different elements. (cation is Hydrogen) Naming: hydro + root element modified “to sound good” + ic Ex: hydrosulfuric acid – H2S Ex: H2Te – hydrotelluric acid Ternary Acids – usually contain three different elements. (cation is Hydrogen; usually anion is a polyatomic ion containing a non-metal and usually the last element is oxygen) Naming: Most Common Form: arises from the “ate” polyatomic ions no prefix + root nonmetal + ic One Less Oxygen Form: arises from the “ite” polyatomic ions no prefix + root nonmetal + ous Two Less Oxygen Form: arises from nothing hypo + root nonmetal + ous One More Oxygen Form: arises from nothing per + root nonmetal + ic One More Oxygen Most Common One Less Oxygen Two Less Oxygen per + root + ic root + ic root + ous hypo + root + ous H(NO4) – pernitric acid H(NO3) – nitric acid H(NO2) – nitrous acid H(NO) – hyponitrous acid H2(SO5) – persulfuric acid H2(SO4) – sulfuric acid H2(SO3) – sulfurous acid H2(SO2) – hyposulfurous H2(CO4) – percarbonic acid H2(CO3) – carbonic acid H2(CO2) – carbonous acid H2(CO) – hypocarbonous H(ClO4) – perchloric acid H(ClO3) – chloric acid H(ClO2) – chlorous acid H(ClO) - hypochlorous *the number of hydrogen stays the same because the oxidation numbers stay the same when changing oxygen* Ex: Name: H2SeO2 H2(SeO4) – selenic acid hyposelenous acid (two less oxygen) Ex: Write: tellurous acid H2(TeO4) – telluric acid H2TeO3 (one less oxygen) Ex: Write: periodic acid H(IO3) – iodic acid H(IO4) (one more oxygen) Ex: Write: HBr hydrobromic acid Acidic Anhydrides – any nonmetal oxide which when dissolved in water will produce an acid (synthesis reaction). Ex: CO2 + H2O  H2(CO3) Ex: SO2 + H2O  H2(SO3) ***ACID RAIN*** Basic Anhydrides – any metallic oxide which when dissolved in water will produce a base (synthesis reaction). Ex: Na2O + H2O  2Na(OH) Neutralization Reactions / Titration Reactions Titration – An experiment involving an acid/a base with a known value and an acid/a base with an unknown value. Acid + Base  Salt + Water The + ion of the acid combines with the – ion of the base to form the salt. Ex: If 22.0 mL of 0.100M HCl (aq) is required to neutralize 31.0 mL NaOH (aq), determine the molarity of the Na(OH) solution. Balanced Equation: HCl + Na(OH)  NaCl + H(OH) 22mL HCl per 31mL NaOH M of HCl Mole-to-Mole Ratio 22.0 mL HCl .100 mol HCl 1 mol NaOH --------------------x ------------------- x ----------------- = 0.071M NaOH 31.0 mL NaOH 1L HCl 1 mol HCl Net Ionic Equations

- written for reactions that take place in H2O and show only the ions that participate in the reaction Spectator Ions – present, but don’t participate in the reaction Considerations 1) Soluble salts written in IONIC FORM. 2) Most other stuff is written in MOLECULAR FORM. 3) Strong acids & bases written in IONIC FORM. 4) Weak acids & bases written in MOLECULAR FORM. Rules 1) Binary Acids: Only three strong (HCl, HBr, HI) and all other are weak acids. 2) Ternary Acids: # of Oxygens - # of Hydrogens = ___ If the answer is 2 more, then the acid is strong. 3) Polyprotic Acids: Acids with more than one ionizable Hydrogen ion. The second acid is ALWAYS weak. Ex: H2(SO4)  H+ + (HSO4)4) Bases: Hydroxides of groups IA and IIA are ALL STRONG. Two exceptions: Be(OH)2 & NH3 5) Salts: Soluble (AQ) – ionic form. Insoluble (CR) – molecular form. 6) Oxides: ALWAYS MOLECULAR FORM 7) Gases: MOLECULAR FORM 8) Water: MOLECULAR FORM Steps 1) Assign oxidation numbers 2) Identify (molecular/ionic form) 3) Do the “math thing” 1+ 1-

2+

1-

2+

1-

3+ 1-

1+

1-

1+ 1-

Ex: 4HCl (aq) + 2Cr(NO3)2 (aq) + 2HgCl2 (aq)  2CrCl3 (aq) + Hg2Cl2 (cr) + 4H(NO3) (aq) strong acid

soluble salt

soluble salt

soluble salt

salt

strong acid

4H1+ + 4Cl1- + 2Cr2+ + 4(NO3)1- + 2Hg2+ + 2Cl1-4Cl1-  2Cr3+ + 6Cl1- + Hg2Cl2 + 4H1+ + 4(NO3)12Cr2+ + 2Hg2+ + 2Cl1-  2Cr3+ + Hg2Cl2

LOWEST

decreases

increases

electronegativity increases

HIGHEST

Why Acids Act Like Acids and Bases Act Like Bases Electronegativity – the ability/tendency of an atom to attract a shared pair of electrons to itself (chemical bonds). Groups IA & IIA – lowest electronegativity (They are trying to get rid of one or two electrons, not gain more.) Groups on Right – highest electronegativity (They only need to attract a few electrons before they are full.) The Shielding Effect – When filled inner energy levels block the attraction of the nucleus from reaching the outer electrons F  highest electronegativity electronegativity decreases (little shielding effect)

lowest electronegativity  Fr 1. distance of electrons from nucleus 2. the shielding effect Most Common Elements in Acids & Bases: H : O : X ; if element X is very HIGHLY electronegative (nonmetal), then the electron pair it shared electrons is sharing with Oxygen will be more attracted by X. In water, X will take the electrons from Oxygen, which will in turn take electrons from Hydrogen. When + ionization occurs, H ions are produced. ACTS LIKE AN ACID. ; if element X is very LOWLY electronegative (metal), then the electron pair will be more attracted by O. In water, O will take the electron pair. When ionization occurs, X+ ions and (OH)- ions are produced. ACTS LIKE A BASE.

Amphoteric – acts as an acid and a base; ex: water. Haber Process – extracting atmospheric nitrogen and making ammonia (NH3) Allotrope – two different forms of the same element; ex: graphite and diamonds Galvanization – coating Iron with Zinc to prevent rusting. Entropy - disorder REDOX Reactions: (change in oxidation numbers) Half-Reaction Method – break reaction into two half-reactions, balance individually, and then combine. Double Displacement reactions are NEVER REDOX. Oxidation-Reduction reactions are the reactions that involve changes in electron structure. OXIDATION can’t occur without REDUCTION and REDUCTION can’t occur without OXIDATION. Oxidation – involves the loss of electrons; ex: rusting (slow) & burning (fast) Reduction – process by which electrons are added to an atom or ion O I L R I G x s o e s a REDuction i s d i d i u n negative positive a n c i t g t n OXidation i i g o o n n RULES **When writing oxidation numbers, include elements in polyatomic ions** 1) Free Elements – ALWAYS ZERO 2) Find monatomic ions using periodic table. 3) Hydrogen is normally 1+ (exception: LiH, H is 1-) 4) Oxygen is 2- (exception: peroxides, O is 1-) 5) Group IA: 1+, IIA: 2+, IIIA: 3+ 6) The sum of the oxidation numbers in a polyatomic ion must equal the charge of the ion. Ex: (SO4)2- 1(X) + 4(2-) = 2X = 6+ **The element being REDUCED is found in the OXIDING agent and vice versa** 1+ 5+ 2-

1+

3+ 2-

2+ 2-

1+

5+ 2-

1+ 2-

Ex: H(NO3) + H3(PO3)  NO + H3(PO4) + H2O N reduced, oxidizing agent: NO3; P oxidized, reducing agent: H3PO3 Half-Reaction Method RULES 1) The electrons lost in oxidation MUST equal the electrons gained in reduction. 2) Since these are acidic reactions, if you need some oxygens to help you balance the equation, plenty of H2Os are available. 3) Since these are acidic reactions, if you need H+ ions, there are plenty in solution, just add them in. Balancing: 1. by electrons 2. by atoms 3. by charge (checking) 1+ 5+ 2-

1+

3+ 2-

2+ 2-

1+

5+ 2-

1+ 2-

Ex: H(NO3) + H3(PO3)  NO + H3(PO4) + H2O strong acid

weak acid

oxide

weak acid

Reduction (4H+ + 3e- + NO3  NO + 2H2O) x 2 4+

3-

5+ 6-

+

2+ 2-

-

4+ 4-

water

Oxidation (H2O + H3(PO3)  H3(PO4) + 2e- + 2H+) x 3 2+ 2-

3+ 3+ 6-

3+ 5+ 8-

2-

-

2+ +

82H + 6e + 2NO3 + 3H2O + 3H3(PO3)  2NO + 4H2O + 3H3(PO4) + 6e + 6H 2H+ + 2NO3 + 3H3PO3  2NO + H2O + 3H3PO4

pH and pOH pH – a measure of hydronium ion concentration – [H3O+] pOH – a measure of hydroxide ion concentration – [OH-] on the pH scale: 0 ≤ acidic < 7 < basic ≤ 14 pH + pOH = 14 [H3O+] x [OH-] = 1 x 10-14 pH = -log[H3O+] pOH = -log[OH-] Ex: If [H3O+] is 10-7, what is the pH? 7 Ex: If the pH is 8, what is [H3O+]? 10-8 Ex: If [H3O+] is 4.37 x 10-4, what is the pH? pH = -log[4.37 x 10-4] = -(log 4.37 + log 10-4) = 3.36 + Ex: If the pH is 2.3, what is [H3O ]? 2.3 = -log[H3O+] -2.3 = log[H3O+] 10-2.3 = [H3O+] = 5 x 10-3 Radioactive Decay alpha – nuclei of He atoms beta – electrons gamma – electromagnetic radiation (most common form of penetrating radiation) alpha + beta = He atom beyond Z=83, all radioactive Inverse Square Law - some physical quantity or strength is inversely proportional to the square of the distance from the source of that physical quantity. (twice as far away, ¼ the effect) Strong Interaction – stronger than the force of repulsion between two protons

mass (A) protons (Z)

238 92

4 2

He

Li

As you add protons, the nucleus gets bigger, and you need more and more neutrons. Transmutation – when atoms change from one form to another. Natural Transmutation – 4 2

234 90

U  He + Th

First Artificial Transmutation –

14 7

5

+ one proton = 3

4 2

1 1

17 8

N + He  H + O

Freezing Point Depression/Boiling Point Elevation Molality – moles of solute per kg (1000g) of H2O Ex: What is the freezing-point depression of water in a solution of 35g of glucose, C6H12O6, and 200.0g of water? (Molal freezing point depression constant for water is -1.86°C/m.) 35g C6H12O6 1 mol C6H12O6 1000g H2O -1.86°C ------------------- x --------------------- x --------------- x ------------ = -1.8°C 200g H2O 180g C6H12O6 1 kg H2O 1 mol/kg Ex: What is the expected boiling-point elevation for a solution that contains 100g of calcium nitrate, Ca(NO3)2, and 700g of water? (Molal boiling point elevation constant for water is 0.51°C/m.) 100g Ca(NO3)2 1 mol Ca(NO3)2 1000g H2O 0.51°C --------------------- x ----------------------- x --------------- x ------------- = .44°C 700g H2O 164.1g Ca(NO3)2 1 kg H2O 1 mol/kg