EXAM 1 REWIEW

Chapter 1: Foundations of Chemistry

1. Structure of the Atom

• Atom – The smallest unit of matter that retains the properties of an element. Atoms consist of a nucleus (containing protons and neutrons) and electrons orbiting the nucleus.

• Proton – A positively charged subatomic particle found in the nucleus.

• Neutron – A neutrally charged subatomic particle in the nucleus.

• Electron – A negatively charged subatomic particle that orbits the nucleus in electron clouds.

• Nucleus – The dense center of an atom, containing protons and neutrons.

2. Scientific Method

The Scientific Method is a systematic approach used in scientific study to explore observations and answer questions. The steps include:

1. Observation – Identifying a problem or question.

2. Hypothesis – A testable statement that explains an observation.

3. Experimentation – A controlled method to test the hypothesis.

4. Analysis – Interpreting data collected from the experiment.

5. Conclusion – Determining whether the hypothesis was correct or not.

3. Theory vs. Law

• Scientific Theory – A well-tested explanation for a broad set of observations. Theories explain why things happen (e.g., Atomic Theory).

• Scientific Law – A statement that describes a pattern in nature but does not explain why it happens (e.g., Law of Conservation of Mass).

4. Classification of Matter

• Matter – Anything that has mass and takes up space.

• Pure Substance – A type of matter with a fixed composition (cannot be separated by physical means).

  • Element – A pure substance that consists of only one type of atom (e.g., Oxygen, Gold).

  • Compound – A pure substance composed of two or more elements chemically bonded together (e.g., Water - H₂O).

• Mixture – A combination of two or more substances that are not chemically bonded.

  • Homogeneous Mixture – A mixture that is uniform in composition (e.g., saltwater, air).

  • Heterogeneous Mixture – A mixture where different substances are visibly distinct (e.g., salad, sand in water).

• Crystalline Solid – A solid with a regular, repeating atomic structure (e.g., salt, diamonds).

• Amorphous Solid – A solid with no fixed atomic structure (e.g., glass, rubber).

5. Methods of Separation

• Distillation – A method that separates substances based on boiling points (e.g., purifying water by boiling and condensing it).

• Decanting – Separating liquid from solid by pouring off the liquid without disturbing the solid.

• Filtration – Using a filter to separate a solid from a liquid (e.g., filtering coffee grounds from brewed coffee).

6. Physical vs. Chemical Properties and Changes

• Physical Property – A characteristic that can be observed without changing the substance’s identity (e.g., color, boiling point, density).

• Chemical Property – A characteristic that describes a substance’s ability to undergo a chemical change (e.g., flammability, reactivity).

• Physical Change – A change that does not alter the substance’s identity (e.g., melting ice, cutting paper).

• Chemical Change – A change that results in a new substance being formed (e.g., burning wood, rusting iron).

7. Energy and Its Forms

• Kinetic Energy – Energy of motion (e.g., a moving car).

• Potential Energy – Stored energy due to position or composition (e.g., a stretched rubber band).

• Thermal Energy – Energy associated with the temperature of a substance (heat).

• Chemical Energy – Energy stored in chemical bonds that is released during reactions (e.g., energy stored in food).

8. Measurement and Units

• SI System – The International System of Units used in science. Common units include meters (m) for length, kilograms (kg) for mass, and seconds (s) for time.

• Metric vs. English System – The metric system (SI) is used in science, while the English system is used in some countries (e.g., feet, pounds).

• Numerical Prefixes –

  • Mega (M) – 1,000,000 (10⁶)

  • Kilo (k) – 1,000 (10³)

  • Deci (d) – 0.1 (10⁻¹)

  • Centi (c) – 0.01 (10⁻²)

  • Milli (m) – 0.001 (10⁻³)

  • Micro (µ) – 0.000001 (10⁻⁶)

9. Temperature Conversions

• Kelvin (K) to Celsius (°C):K=°C+273.15K=°C+273.15

• Celsius (°C) to Fahrenheit (°F):°F=(°C×95)+32°F=(°C×59​)+32

• Fahrenheit (°F) to Celsius (°C):°C=(°F−32)×59°C=(°F−32)×95​

10. Density

Density is the mass per unit volume of a substance.

Density=MassVolumeDensity=VolumeMass​

• Measured in g/mL or g/cm³.

11. Significant Figures

Rules for Multiplication/Division:

• The result should have the same number of significant figures as the number with the fewest significant figures in the calculation.

Rules for Addition/Subtraction:

• The result should be rounded to the least precise decimal place.

12. Precision vs. Accuracy

• Precision – How close multiple measurements are to each other.

• Accuracy – How close a measurement is to the actual value.

13. Conversion Factors and Dimensional Analysis

• Dimensional Analysis – A method for converting units by multiplying by conversion factors.
  Example: Convert 5.0 km to meters

5.0 km×1000 m1 km=5000 m5.0 km×1 km1000 m​=5000 m

14. Algebra in Chemistry

Being able to rearrange equations to solve for a desired variable is critical.
Example: Solve for V in the density equation:

D=M/V

Rearrange:

V=M/D

Chapter 2: Atomic Theory and the Periodic Table

1. Scanning Tunneling Microscope (STM)

• A specialized instrument used to visualize individual atoms by detecting the electrical current between a fine tip and the surface of a material.

• Helped confirm the atomic structure and revolutionized nanotechnology.

2. Laws Governing Atomic Theory

Law of Conservation of Mass

• Mass is neither created nor destroyed in a chemical reaction.

• This means the total mass of the reactants equals the total mass of the products.

• Example: If you burn 10 g of wood and get 2 g of ash, the missing 8 g is in the form of gases like CO₂ and H₂O vapor.

Law of Definite Proportions

• A given compound always contains the same elements in the same proportions by mass.

• Example: Water (H₂O) always consists of 88.8% oxygen and 11.2% hydrogen by mass, no matter the sample.

Law of Multiple Proportions

• When two elements form more than one compound, the mass ratio of one element to a fixed mass of the other is always a whole number.

• Example: CO and CO₂ – Carbon forms two different compounds with oxygen. The ratio of oxygen in CO₂ to oxygen in CO is 2:1, a simple whole number ratio.

3. Dalton’s Atomic Theory

John Dalton proposed the first modern atomic theory in 1808. His theory states:

1. All matter is made of tiny, indivisible particles called atoms.

2. Atoms of the same element are identical, but different from atoms of other elements.

3. Atoms combine in whole-number ratios to form compounds.

4. Chemical reactions involve rearrangement of atoms, but atoms themselves are not created or destroyed.

Some parts of Dalton’s theory were later revised (e.g., atoms are divisible into subatomic particles, and isotopes exist), but it was a foundation for modern chemistry.

4. Discovery of Subatomic Particles

J.J. Thomson and the Electron

• Discovered the electron in 1897 using the cathode ray tube experiment.

• Proposed the "plum pudding model", where electrons were scattered inside a positively charged "pudding."

Millikan’s Oil Drop Experiment

• Determined the charge of an electron to be −1.60 × 10⁻¹⁹ C and calculated the electron's mass.

Rutherford’s Gold Foil Experiment

• Showed that atoms have a small, dense, positively charged nucleus instead of the "plum pudding" model.

• Led to the nuclear model of the atom.

5. Atomic Structure and Subatomic Particles

6. Isotopes

• Isotopes are atoms of the same element with different numbers of neutrons.

• Example: Carbon has three isotopes – C-12, C-13, C-14.

• Isotopes are written as A/Z X (e.g., ¹²₆C), where:

  • A = Mass Number (protons + neutrons)

  • Z = Atomic Number (protons)

7. Periodic Table and Periodic Trends

Metals, Nonmetals, and Metalloids

• Metals – Conduct electricity, malleable, ductile, and shiny (e.g., Fe, Cu).

• Nonmetals – Poor conductors, brittle, and can be gases or solids (e.g., O, S).

• Metalloids – Have properties of both metals and nonmetals (e.g., Si, B).

Periodic Trends

• Alkali Metals (Group 1) – Highly reactive, soft metals (e.g., Na, K).

• Alkaline Earth Metals (Group 2) – Less reactive than alkali metals but still form strong bases (e.g., Mg, Ca).

• Halogens (Group 17) – Highly reactive nonmetals that form salts with metals (e.g., Cl, F).

• Noble Gases (Group 18) – Inert gases with full valence shells (e.g., He, Ne).

8. Atomic Mass and Molar Mass

• Atomic Mass – The weighted average mass of an element’s isotopes (measured in amu).

• Molar Mass – The mass of one mole of a substance (g/mol).

Calculating Weighted Atomic Mass

Average Atomic Mass=(% abundance×mass)+(% abundance×mass)+...Average Atomic Mass=(% abundance×mass)+(% abundance×mass)+...

9. Avogadro’s Number and the Mole

• Avogadro’s Number (6.022 × 10²³ particles/mol) – The number of atoms, ions, or molecules in one mole of a substance.

• Used to convert between atoms/molecules and moles.

Example: Convert 3.00 moles of oxygen atoms to atoms:

3.00 mol×6.022×1023 atoms1 mol=1.81×1024 atoms3.00 mol×1 mol6.022×1023 atoms​=1.81×1024 atoms

10. Mass Spectrometry

• A technique used to determine the mass of atoms and molecules.

• Helps identify isotopes and calculate atomic mass

Chapter 3: Chemical Compounds and Bonding

1. Elements vs. Compounds

• Element – A pure substance made of only one type of atom (e.g., O₂, Fe).

• Compound – A substance composed of two or more different elements chemically bonded in a fixed ratio (e.g., H₂O, CO₂).

• Key idea: Compounds have properties different from the individual elements that make them up.

Example:

• Sodium (Na) is a highly reactive metal, and chlorine (Cl) is a poisonous gas.

• But together, they form sodium chloride (NaCl), which is table salt!

2. Chemical Bonds

Ionic Bonds

• Formed between a metal and a nonmetal due to the transfer of electrons.

• Metals lose electrons and become cations (+).

• Nonmetals gain electrons and become anions (−).

• Example: NaCl → Na⁺ + Cl⁻

Covalent Bonds

• Formed between two nonmetals by sharing electrons.

• No ions are formed.

• Example: H₂O → Oxygen shares electrons with hydrogen atoms.

3. Identifying Ionic vs. Molecular (Covalent) Compounds

Example:

• NaCl (ionic) – metal (Na) + nonmetal (Cl).

• CO₂ (molecular) – nonmetal (C) + nonmetal (O).

4. Molecular vs. Empirical Formulas

• Molecular Formula – Shows the actual number of atoms (e.g., H₂O₂).

• Empirical Formula – Shows the simplest ratio (e.g., HO for H₂O₂).

Example:

• Glucose = C₆H₁₂O₆ (molecular formula)

• Empirical formula = CH₂O (simplest ratio)

5. Molecular Elements

• Some elements exist as diatomic molecules instead of single atoms.

• Mnemonic: "BrINClHOF" (Br₂, I₂, N₂, Cl₂, H₂, O₂, F₂)

• Example: Oxygen exists as O₂, not just O.

6. Naming Chemical Compounds

A. Naming Ionic Compounds

Rule: Name the metal first, then the nonmetal with “-ide” ending.

• NaCl = Sodium chloride

• CaO = Calcium oxide

Transition Metals (Variable Charge)

• Use Roman numerals to indicate the charge.

• FeCl₂ = Iron (II) chloride

• FeCl₃ = Iron (III) chloride

B. Naming Molecular (Covalent) Compounds

Rule: Use prefixes to indicate the number of atoms.

Examples:

• CO₂ = Carbon dioxide

• PCl₅ = Phosphorus pentachloride

Note: "Mono-" is not used for the first element (e.g., CO is carbon monoxide, not monocarbon monoxide).

C. Naming Acids

7. Polyatomic Ions

• Ions composed of multiple atoms bonded together.

• Common Polyatomic Ions:

Trends:

• Per- = 1 more oxygen (e.g., ClO₄⁻ = Perchlorate)

• -ate = Base form (e.g., ClO₃⁻ = Chlorate)

• -ite = 1 less oxygen (e.g., ClO₂⁻ = Chlorite)

• Hypo- = 2 less oxygen (e.g., ClO⁻ = Hypochlorite)

8. Formula Mass and Molar Mass

• Formula Mass – The sum of atomic masses of all atoms in a compound (amu).

• Molar Mass – Mass of 1 mole of a compound (g/mol).

Example (H₂O):

Formula Mass=(2×1.008)+(1×16.00)=18.016 amuFormula Mass=(2×1.008)+(1×16.00)=18.016 amu

9. Mole Conversions

Use Avogadro’s number (6.022 × 10²³ particles/mol) for conversions.

Example: Convert 10 g of NaCl to moles.

10 g÷58.44 g/mol=0.171 mol10 g÷58.44 g/mol=0.171 mol

10. Empirical and Molecular Formulas

Finding Empirical Formula

1. Convert percent composition to moles.

2. Divide by the smallest number of moles.

3. Multiply to get whole numbers if necessary.

Finding Molecular Formula

1. Find empirical formula mass.

2. Divide molecular molar mass by empirical mass.

3. Multiply the empirical formula by that number.

Example:

• Empirical formula: CH₂O

• Molecular mass: 180 g/mol

• Empirical mass: 30 g/mol

• Molecular formula = C₆H₁₂O₆ (since 180 ÷ 30 = 6).

11. Balancing Chemical Equations

• The Law of Conservation of Mass requires equal atoms of each element on both sides.

• Use coefficients, not subscripts, to balance.

Example:

Unbalanced: H2+O2→H2OUnbalanced: H2​+O2​→H2​OBalanced: 2H2+O2→2H2OBalanced: 2H2​+O2​→2H2​O

Chapter 4: Reaction Stoichiometry

1. Balancing Chemical Equations

• Why? The Law of Conservation of Mass states that matter cannot be created or destroyed. So, the number of atoms must be equal on both sides of a reaction.

• How? Adjust coefficients (big numbers before compounds), not subscripts (small numbers in formulas).

Example:

Unbalanced: H2+O2→H2OUnbalanced: H2​+O2​→H2​OBalanced: 2H2+O2→2H2OBalanced: 2H2​+O2​→2H2​O

Now, there are 4 hydrogen atoms and 2 oxygen atoms on both sides.

2. Using Mole Ratios from a Balanced Chemical Equation

• Definition: The mole ratio tells how many moles of one substance react with another.

• Found in: The coefficients of a balanced equation.

Example:

2H2+O2→2H2O2H2​+O2​→2H2​O

Mole ratios:

• 2 moles H₂ react with 1 mole O₂

• 2 moles H₂O are produced per 2 moles H₂

3. Limiting Reactant (LR) & Excess Reactant (ER)

• Limiting Reactant (LR): The reactant that runs out first and limits the amount of product formed.

• Excess Reactant (ER): The reactant that is left over after the reaction is complete.

Steps to Find LR:

1. Convert given grams to moles.

2. Use the mole ratio from the balanced equation.

3. Determine which reactant produces the least amount of product.

Example:

2H2+O2→2H2O2H2​+O2​→2H2​O

• If given 10 moles H₂ and 4 moles O₂:

  • 10 moles H₂ produces 10 moles H₂O (1:1 ratio).

  • 4 moles O₂ produces 8 moles H₂O (1:2 ratio).

  • O₂ is the LR because it produces the least amount of product.

4. Theoretical Yield & Percent Yield

• Theoretical Yield: The maximum amount of product that can be formed (calculated from LR).

• Actual Yield: The amount of product actually obtained from the experiment.

• Percent Yield: Measures efficiency of a reaction.

%Yield=(Actual YieldTheoretical Yield)×100%Yield=(Theoretical YieldActual Yield​)×100

Example:

• Theoretical yield = 20 g H₂O

• Actual yield = 15 g H₂O

%Yield=(1520)×100=75%%Yield=(2015​)×100=75%

5. Molarity (MM)

• Definition: Concentration of a solution, measured in moles of solute per liter of solution.

M=moles of soluteliters of solutionM=liters of solutionmoles of solute​

Example:

• 0.5 moles of NaCl in 2 L solution

M=0.5 moles2 L=0.25MM=2 L0.5 moles​=0.25M

6. Dilution Formula

• Used to dilute a solution (reduce its concentration).

M1V1=M2V2M1​V1​=M2​V2​

Where:

• M1,V1M1​,V1​ = Initial molarity & volume

• M2,V2M2​,V2​ = Final molarity & volume

Example:

• You have 3.0 M HCl and want 500 mL of 1.5 M HCl.

• How much stock solution (V1V1​) do you need?

(3.0M)(V1)=(1.5M)(0.500L)(3.0M)(V1​)=(1.5M)(0.500L)V1=(1.5M)(0.500L)3.0M=0.250L=250 mLV1​=3.0M(1.5M)(0.500L)​=0.250L=250 mL

So, you need 250 mL of 3.0 M HCl and dilute it to 500 mL.

7. Types of Reactions

A. Acid-Base Reactions

• Acid: Donates H⁺ (proton).

• Base: Accepts H⁺.

• Neutralization Reaction: Acid + Base → Water + Salt.

Example:

HCl+NaOH→NaCl+H2OHCl+NaOH→NaCl+H2​O

B. Gas Evolution Reactions

• Gas is produced as a product.

• Example: Carbonates (CO32−CO32−​) react with acids to form CO2CO2​ gas.

Na2CO3+2HCl→CO2+H2O+2NaClNa2​CO3​+2HCl→CO2​+H2​O+2NaCl

C. Precipitation Reactions

• Two aqueous solutions form a solid precipitate.

• Use a solubility table to determine if a precipitate forms.

Example:

AgNO3+NaCl→AgCl(s)+NaNO3AgNO3​+NaCl→AgCl(s)+NaNO3​

AgCl is insoluble → precipitate forms!

D. Oxidation-Reduction (Redox) Reactions

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

Example:

Zn+CuSO4→ZnSO4+CuZn+CuSO4​→ZnSO4​+Cu

• Zn is oxidized (loses electrons).

• Cu²⁺ is reduced (gains electrons).

8. Oxidation States & Redox Agents

• Oxidation State: A number assigned to an element based on electron transfer.

• Rules for Assigning Oxidation Numbers:

  1. Free elements (O₂, N₂) = 0.

  2. Group 1 metals = +1, Group 2 = +2.

  3. Oxygen = -2 (except in peroxides, where it’s -1).

  4. Hydrogen = +1 (except in metal hydrides, where it’s -1).

Example: Find oxidation states in H2OH2​O.

• H = +1

• O = -2

9. Identifying Spectator Ions

• Spectator Ions: Ions that do not change in a reaction.

• Net Ionic Equation: Only includes reacting species.

Example:

Molecular: AgNO3+NaCl→AgCl(s)+NaNO3Molecular: AgNO3​+NaCl→AgCl(s)+NaNO3​

Complete Ionic:

Ag++NO3−+Na++Cl−→AgCl(s)+Na++NO3−Ag++NO3−​+Na++Cl−→AgCl(s)+Na++NO3−​

Net Ionic Equation:

Ag++Cl−→AgCl(s)Ag++Cl−→AgCl(s)

Spectator Ions: Na⁺, NO₃⁻ (they remain unchanged).