Periodic Table of Elements

Sourvelis' Problem Solving Strategies

How To FRQ

Metric Unit Conversions

Chemistry Symbol Guide

Common Ion Guide

Equations and Constants

Orbital Filling Order

Electronegativity Values

VSEPR Theory Table

Perdioic Table of Ions

Nomenclature Flowchart

Chemical Reaction Rules

Activity Series of Metals

Solubility Guidelines

Molar Conversion - Stoichiometry

This is a guide to how to answer open-ended response questions. A well-answered free response question will demonstrate mastery of the AP Chemistry science practices and contain elements of the following:

Math and Data Analysis

1. Show Every Step: Always write out your dimensional analysis, formulas used, and units throughout. Label quantities clearly. Use proper mathematical routines as defined by Practice 5.

2. Use Units and Significant Figures: Units must be included in your final answer, and proper sig figs should reflect the given data. You may lose credit for incorrect or missing units (Practice 5.F).

3. Keep Going: If you make a mistake early on but apply it correctly in the rest of your work, you can still earn follow-through credit. Don’t stop!

Written Justification and Scientific Argumentation

1. Answer with Evidence + Reasoning: Your response should always support claims using chemical principles, laws, or data. Use Practice 6 to guide this (e.g., “because of Coulomb’s Law…”).

2. Be Specific: Refer directly to the substance or concept. Avoid vague pronouns. For instance, say “the K⁺ ion” not “it.”

3. Use Proper Terminology: Use precise chemical vocabulary (e.g., ionization energy, molar enthalpy, rate law) from the CED. This is essential for earning credit.

4. Be Concise and Coherent: Keep your explanations focused. Avoid rambling, as contradictory or irrelevant information can lead to point deductions.

5. Model, Graph, or Diagram: If asked, or if it helps clarify your explanation, include a diagram, Lewis structure, or particle model. Practice 3 and 4 emphasize data and model representation.

AP Chemistry Science Practices You’ll Be Scored On:

• Practice 1: Models and Representations
• Practice 2: Question and Method
• Practice 3: Representing Data and Phenomena
• Practice 4: Model Analysis
• Practice 5: Mathematical Routines
• Practice 6: Argumentation

As you practice FRQs, review the scoring guidelines and rubrics found on AP Central to ensure your responses align with the expectations.

Algebra can be simplified into three fundamental rules

1. Make sure unknown variable is in the numerator (if it is a fraction).

  - Write down the givens from the problem. Include units.

2. Isolate the unknown variable by moving it or other terms using the opposite mathematical operation that is being applied to it.

  - Ex: the opposite of multipling is dividing.

  - **Note - Sometimes it is better to move the unknown instead of moving everything else.

3. Whatever operation you do to one side of the equation, you must do to all terms on the other side as well.

Givens   - Write down the givens from the problem. Include units.

Unknowns   - Write down the unknow from the problem. Include units.

Equation   - Write down the equation. Determine from above.

Solve   - Solve the literal Equation for the unknown.

Substitution   - Substitute the known varibles and calculate

Part 1: Organize the problem

1. Identify/label the \(given\) and the conversion factor.

2. Identify/label the \(goal\) and the conversion factor.

3. What can be ignored and why?

4. Identify the equation type. __________________

5. Predict the products and write a balanced chemical equation.

Part 2: Calculation

1. Start the problem with the \(given\) and multiply by the conversion factor to calculate moles.

2. Mulitply by the mole-to-mole ratio (\(\frac{goal}{given}\)) from the balanced equation.

3. Multiple by the conversion factor to convert into the desired units for the \(goal\).

Model Problem 1:

What volume of hydrogen gas can be produced when a 0.0243 g sample of magnesium (Mg) is reacted with excess aqueous 6 M hydrochloric acid (HCl)?

Solution:

Elemental Names - assigned by the discoverer of that element.

• origin and language varies wildly.
• named after planets: mercury, uranium, neptunium, plutonium
• named after properties: bromine and osmium, mean stench and smells.

Elemental Symbols - varied wildly and weirdly. Many element symbols were abbreviated on the basis of their Greek or Latin name.

• Sodium - \(\text{Na}\) from the Latin word Natrium:
• Tungsten - \(\text{W}\) from the German word Wolfran

Elements are represented by one or two-letter symbols, where the first letter is capitalized, and the second (if present) is lowercase.

• One letter: \(\text{H}\) – Hydrogen    \(\text{O}\) – Oxygen
• Two letters: \(\text{Na}\) – Sodium    \(\text{Cl}\) – Chlorine

A polyatomic ion is an ion composed of two or more atoms that are covalently bonded together, acting as a single charged entity. The whole group of atoms has a net positive or negative charge, which allows it to participate in ionic bonds with other ions.
Memorizing the names, formulas, and charges of common polyatomic ions is a requirement.
  Memorizing polyatomic ions makes problem-solving faster and builds a solid foundation for tackling complex chemistry topics

Ionic Charge – The actual electrical charge on an ion, resulting from the loss or gain of electrons. For main-group elements, the charge often corresponds to how many electrons an atom must lose or gain to achieve a stable octet (eight valence electrons), similar to a noble gas electron configuration. When this is achieved, we say the ion is isoelectronic.
Chlorine has 7 valence electrons. If chlorine gains one, it will have 17 \(p^+\) and 18 \(e^-\), resulting in a negative charge overall. Now chlorine will have 3 principal energy levels that are completely full—just like (isoelectronic with) Argon.

Cations are positively charged ions formed when atoms lose one or more electrons. Metals, especially those in Groups 1, 2, and 13, commonly form cations.

Anions are negatively charged ions formed when atoms gain one or more electrons. Nonmetals in Groups 15, 16, and 17 typically form anions.

Sodium \(\text{Na}\) is a metal that has 11 protons and 11 electrons. In its atomic state, it has a neutral charge of \(0\). \(\text{Na}\) is highly reactive, however. In fact, you can't find \(\text{Na}\) in nature in its atomic form. It readily bonds with chlorine \(\text{Cl}\) by giving one valence electron to chlorine. In doing this, both atoms achieve a stable electron arrangement—the octet referred to earlier. \(\text{Na}\) becomes isoelectronic with the noble gas \(\text{Ne}\), and \(\text{Cl}\) with \(\text{Ar}\).
\(\text{Cl}\) gains one electron \([\text{Cl}^-]\) and becomes an anion with a –1 charge. The magnitude is 1, and the charge is negative. \([\text{Cl}^-]\) is called chloride.
\(\text{Na}\) loses one electron \([\text{Na}^+]\) and becomes a cation with a +1 charge. The magnitude is 1, and the charge is positive. \([\text{Na}^+]\) is still called sodium. Only the name of anions changes by switching the suffix to "ide".

Avogadro's number, \( N_A = 6.022 \times 10^{23} \, \text{mol}^{-1} \)

Speed of light, \( c = 2.998 \times 10^8 \, \text{m/s} \)

Planck's constant, \( h = 6.626 \times 10^{-34} \, \text{J} \cdot \text{s} \)

1 atm = 760 torr = 760 mm Hg = 101,325 Pa = 101.325 kPa = 29.9 in. Hg = 14.7 psi

Instructions - When answering questions about electronic geometry and molecular geometry use the following table. First, count the number of atoms and lone pairs stuck to the "central" atom. Find that number in the first column. The total number of atoms and lone pairs corresponds to the Electron-pair Geometry. Lone pairs of electrons repel bonding pairs and change the shape of the molecule. We call that molecular geometry. Use the table to determine the number of lone pairs and the molecular geometry.

1. All common compounds of Group I (alkali metals) and ammonium ions are soluble.

2. All nitrates, acetates, and chlorates are soluble.

3. All binary compounds of the halogens (other than F) with metals are soluble, except those of Ag, Hg(I), and Pb. (Pb halides are soluble in hot water.)

4. All sulfates are soluble, except those of barium, strontium, calcium, lead, silver, and mercury (I). The latter three are slightly soluble.

5. Except for rule 1, carbonates, hydroxides, oxides, silicates, and phosphates are insoluble.

6. Sulfides are insoluble except for calcium, barium, strontium, magnesium, sodium, potassium, and ammonium.

1. All common compounds of Group I (alkali metals) and ammonium ions are soluble.

2. All nitrates, acetates, and chlorates are soluble.

3. All binary compounds of the halogens (other than F) with metals are soluble, except those of Ag, Hg(I), and Pb. (Pb halides are soluble in hot water.)

4. All sulfates are soluble, except those of barium, strontium, calcium, lead, silver, and mercury (I). The latter three are slightly soluble.

5. Except for rule 1, carbonates, hydroxides, oxides, silicates, and phosphates are insoluble.

6. Sulfides are insoluble except for calcium, barium, strontium, magnesium, sodium, potassium, and ammonium.

0. Stoichiometry can be used in determining amount of substance reacting or formed.

1. Balance the equation.

2. Decide what quantity of each substance is given. Write it below the substance (one line down if a mole quantity, two lines down if a laboratory quantity).

3. Decide what quantity you want to calculate. Represent it with a question mark directly under the substance (one line down if a mole quantity, two lines down if a laboratory quantity).

4. Complete the diagram with arrows by going from the known laboratory quantity to moles, then to moles of unknown quantity, and finally to laboratory quantity of the unknown. (Step(s) can be omitted if either or both the given quantity and the unknown quantity are moles.)

5. Calculate each quantity following the steps in the diagram.