Chemistry Honors - Units 1-12

Unit 2:

Scientists:

• Dalton: re-proposed the atom theory

  • Stated that mass is neither created or destroyed (law of conservation of mass discovered by Antoine Lavosier)

  • Compound always contains the same elements in the same proportions by mass (law of constant proportions or Proust’s Law)

  • When two elements can form multiple compounds the ratios of masses will remain constant for each compound (Law of multiple proportions by John Dalton)

• JJ Thomson: discovered the electron through the cathode ray tube

  • found that atoms are not indivisible

  • Proposed the Plum Pudding model

• Rutherford discovered the nucleus and the proton

  • used the gold foil experiment

  • Proposed planetary model

• Bohr: proposed the Bohr model of the atom

• Chadwick: discovered the neutron

Subatomic Particles

• Protons

  • in the nucleus

  • +1

  • has mass

• Electron

  • in energy shells

  • -1

  • no mass

• Neutron

  • in the nucleus

  • 0 charge

  • has mass

Mass relationships in atoms:

• Atomic number: the number of protons

• Mass number: the number of protons + neutrons

• How to find the number of electrons: number of protons = number of electrons

Isotopes: atoms with the same atomic number but different mass numbers

• same number of protons

• different number of neutrons

• isotopes are named by their mass number

Beta, Alpha, etc

1. Alpha emission (\frac42He )

   1. Makes you go two left on the periodic table and subtract 4 from the mass number

2. Beta emission (\frac{0}{-1}\beta )

   1. too many neutrons (above the belt of stability)

   2. makes you go right one

3. Positron emission (\frac{0}{+1}\beta )

   1. too many protons (below the belt of stability)

   2. makes you go left one

4. Gamma

   1. Lowers from excited state

Unit 4

Empirical Fomula:

1. If it’s a percentage, assume it’s in grams and bring them all to the mol.

2. divide by the smallest mol

3. put together the equation

% composition and % by mass (same thing)

Percent composition = grams of \frac{element}{compound}\cdot100

Unit 5

Reaction Types:

• Synthesis: A + X → AX

  • If it’s metal oxide + water → metal hydroxide

    • ex:Na_2O+H_2O\to2NaOH

• Decomposition: AX → A + X

  • Metal chlorate → metal chloride + oxygen (O2)

• Double Displacement: AX + BY → AY + BX

• Single Displacement: AX + B → BX + A

• Combustion: C_{x}H_{y}O_{z}+O_2\to CO_2+H_2O

Solubility Rules:

1. All alkali metal compounds are soluble

2. ALl NH4+ compounds are soluble

3. All compounds containing NO3-, ClO3-, ClO4- are soluble

4. Most OH- compounds are insoluble. Alkali metal hydroxides and Ba(OH)2 are exceptions.

5. Most Cl-, Br-, and I- compounds are soluble. The exceptions are: Ag+, Hg2 2+, and Pb 2+

6. All CO3 2-, PO4 3-, and S 2- compounds are insoluble. Exceptions are alkali metal and ammonium compounds.

7. Most SO4 2- compounds are soluble. Exceptions are Ba 2+, Hg 2+, and Pb 2+

• Soluble = aqueous

• Insoluble = precipitate

Unit 6:

• KNOW HOW TO DO ELECTRON CONFIGURATIONS

• Going across to the left and down the periodic table the larger the atom is

  • Positive ions are smaller than regular atoms

  • Negative ions are larger than regular atoms

Unit 7:

Bonding:

• Ionic Bonding: the transfer of electrons between a metal and non-metal

• Covalent Bonding: the sharing of electrons between non-metals

  • As number of bonds increases, the bond length decreases

  • As number of bonds increases, the bond energy increases

Electronegativity:

Bond polarity:

• Non polar covalent: the electronegativity difference is 0

• Polar covalent: the electronegativity difference is < 2.0, and polarity is pulled towards the molecule with higher electronegativity

• Ionic: the electronegativity difference is >= 2.0

Molecular Geometry

Hybridization

Unit 8

Gas Laws

Boyle’s Law: P_1V_1=P_2V_2

Gay-Lussac’s Law: \frac{P_1}{T_1}=\frac{P_2}{T_2}

Charles Law: \frac{V_1}{T_1}=\frac{V_2}{T_2}

WHAT”S VAPOR PRESSURE AND DIFFUSION RATE????

Diffusion Rate: rate2/rate1 = sq root of M2/sq root of M1

Unit 9

Intermolecular forces

• Ion - Ion: strong force between + and - ions

• Ion - Dipole: attractive force between ion and polar molecule

• Dipole - Dipole: attractive force between polar molecules

• Hydrogen bond: strong attractive force between hydrogen of one molecule and a highly electronegative atom of another

  • Strong type of dipole-dipole

• Dispersion: short-range attractive force of all molecules

Strength of each force:

Dispersion < Dipole-Dipole < Hydrogen Bonding < Ion-Dipole < Ion-Ion

Viscosity

Viscosity: the resistance to flow.

• the more cohesion of the molecules, the greater the viscosity

• as temperature increases, viscosity decreases

Warming and Cooling Curves

Q=mc\Delta T is used to raise the temperature of something up to the point that it phase changes (temperature change)

• m = grams

• T = temperature in celsius

Q=ml is used for phase changing

Quantitative Concentrations of Solutions

1. Molarity = M = mol solute/L solution = mol/L

2. Molality = m = mol solute/kg solvent = mol/kg

3. Normality = N = # equivalents/L solution (often with acids and bases)

   1. Multiply molarity by number of H+ for acids and molarity by number of OH- for bases to get normality

Solubility Rules:

• Soluble = if a solution can have a concentration of 0.1 M or more

• Insoluble = if less than 0.1 M

Unit 10

Exothermic and Endothermic

Exothermic: creates bonds and releases energy

Endothermic: breaks bonds and absorbs energy

Enthalpy

Enthalpy (\Delta H ): the amount of heat released or absorbed in a chemical reaction

Exothermic RXN: Enthalpy is negative

Endothermic RXN: Enthalpy is positive

How to calculate Enthalpy:

1. Measure with an experiment

2. Calculate using Heats of Formation data

3. Calculate using average bond energies

4. Calculate using Hess’ Law

Calculating Enthalpy with heats of formation data:

\Delta H=Products-\operatorname{Re}ac\tan ts

• see examples

Calculating Enthalpy with average bond energies:

\Delta H = reactants - products

• see examples

Calculations of Heat (Q)

Q=mc\Delta T is used to raise the temperature of something up to the point that it phase changes (temperature change)

• m = grams

• T = temperature in celsius

IMPORTANT: Any gas at STP: 1 mol = 22.4 L (STP = 273 K and 1 atm)

Entropy

Entropy: a measure of randomness or disorder in a system

• in nature everything proceeds toward maximum entropy and minimum enthalpy

• Entropy increases going from solid → liquid → gas and is positive

• it’s negative if it goes the opposite direction

Equation: \Delta S = products - reactants

• Same thing as enthalpy with average bond energies but reversed

• + = increasing

• - = decreasing

Spontaneous Process

Gibb’s Free Energy (G):

If \Delta G is positive, it’s NOT spontaneous.

If \Delta G is negative, it IS spontaneous.

Equation: \Delta G=\Delta H-T\Delta S

• T = temperature (K)

• Convert entropy from J/K to kJ/K

• Practice!!!

Unit 11

Rate Law:

Rate of reaction equation: =k\left\lbrack A\right\rbrack^{x}\left\lbrack B\right\rbrack^{y}\left\lbrack C\right\rbrack^{z}

• k = rate constant

• x = order of reaction with respect to A

• y = order of reaction with respect to B

• x + y (adding all your exponents) = overall order of reaction

To find the order of a reaction:

\frac{rate_2}{rate_1}=\left\lbrack\frac{\left\lbrack A\right\rbrack_2}{\left\lbrack A\right\rbrack_1}\right\rbrack^{x}

and order (x) = 0, 1, 2, 3

1=2^x → x = 0, etc etc

PRACTICE

Le Chatelier’s Principle

Definition: When a system at equilibrium is disturbed by applying a stress, and new equilibrium position is attained to relieve the stress

1. Temperature’s effects on equilibrium:

   1. Raising temperature (adding heat so endo):

      1. equilibrium is shifted left to reactants

      2. to relieve stress r must decrease and p must increase

   2. Lowering temperature (removing heat so exo):

      1. equilibrium is shifted to the right to products

      2. To relieve stress, p must decrease and r must increase

2. Pressure’s effects on equilibrium

   1. increasing pressure with 2NO_2\to N_2O_4

      1. equilibrium is shifted to the side with fewer gas molecules

         1. So it’s going towards N2O4 since there’s only one, while there’s two of NO2

   2. Decreasing pressure:

      1. Equilibrium is shifted to the side with a greater number of gas molecules

         1. So it’s going towards NO2

3. Concentration’s effects on equilibrium:

   1. SKIPPED COME BACK TO IT

PRACTICE QUESTIONS WITH ICE ON PG 16

Equilibrium Expression

Equilibrium Expression: How equilibrium is measured

K_eq = Equilibrium constant (no units) is CONSTANT at a given temperature

Formula: K_{eq}=\frac{\left\lbrack products\right\rbrack}{\left\lbrack reactants\right\rbrack}

K_c (concentration): used for solutions (aq) or gases

Formula: K_{c}=\frac{\left\lbrack C\right\rbrack^{c}\cdot\left\lbrack D\right\rbrack^{d}}{\left\lbrack A\right\rbrack^{a}\cdot\left\lbrack B\right\rbrack^{b}} with the coefficients as the powers

• Don’t include anything that’s not aqueous or a gas

• Ex: Don’t put in H2O since it’s a liquid

K_p (Pressure): used for gases only

Formula: K_{p}=\frac{P_{C}^{c}\cdot P_{D}^{d}}{P_{A}^{a}\cdot P_{B}^{b}} with the coefficients as the powers

• Don’t include anything that’s not a gas

Formula for Kp and Kc: K_{p}=K_{c}\left(RT\right)^{\Delta n}

• \Delta n = # mol gas products - # mol gas reactants

• T = Temp (K)

• R = 0.0821

KNOW HOW TO DO ICE OR CALCULATE EQUILIBRIUM CONCENTRATIONS FROM INITIAL CONCENTRATIONS ON Pg 12

MAYBE COME BACK FOR REACTION QUOTIENT

Unit 12

pH and pOH

In pure water, both [H+] and [OH-] = 1.0 × 10^-7 M

Formulas for calculating pH and pOH: pH = -log[H+], pOH = -log[OH-] (both are provided)

• pH + pOH = 14 (also provided

K_a and K_b

K_a = acid dissociation constant

Formula: K_c = K_a = \frac{\left\lbrack products\right\rbrack}{\left\lbrack reactants]\right\rbrack} with coefficients as exponents

K_b = base - dissociation constant

Formula is the same thing as K, just make sure to exclude water since it’s not aqueous or a gas

PRACTICE THE QUESTIONS ON PAGES 4 and 5

IMPORTANT: K_{a}\cdot K_{b}=K_{w} (provided)

• They’re acid-base conjugate pairs

• As Ka increases, Kb decreases (inverse relationship)

• The stronger the acid, the weaker it’s conjugate base