CHEMISTRY LAB
"Chemistry is the study of matter, but I prefer to see it as the study of change."
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Key Topics
Atomic Structure
Protons, neutrons, electrons and quantum mechanics
Chemical Bonds
Ionic, covalent, and metallic bonding
Reactions
Balancing equations and reaction types
Organic Chemistry
Hydrocarbons, functional groups, and reactions
Chemistry Notes
Chapter: Periodic Table Trends
📘 Chapter: Periodic Table Trends
1. Introduction
The periodic table arranges elements in order of increasing atomic number. Elements in the same group have similar chemical properties due to similar valence electron configurations. Various periodic trends are observed across periods and down groups.
📘 Chapter: Periodic Table Trends
2. Periodic Trends Visual Guide
| Property | Across a Period → | Down a Group ↓ |
|---|---|---|
| Atomic Radius | ⬇ Decreases | ⬆ Increases |
| Ionization Energy | ⬆ Increases | ⬇ Decreases |
| Electron Affinity | ⬆ More Negative | ⬇ Less Negative |
| Electronegativity | ⬆ Increases | ⬇ Decreases |
| Metallic Character | ⬇ Decreases | ⬆ Increases |
| Non-metallic Character | ⬆ Increases | ⬇ Decreases |
3. Quick Reference Example
- Atomic radius: Na > Cl > F
- Ionization energy: F > Cl > Na
- Electronegativity: F > Cl > Na
4. Refer to Full Periodic Table
4. Refer to Full Periodic Table
For complete details of elements, their symbols, atomic numbers, and properties, visit your full periodic table.
- Groups: Columns → same valence electrons, similar reactivity.
- Periods: Rows → increasing energy levels.
- Trends:
- Atomic radius ↓ across a period, ↑ down a group.
- Ionization energy ↑ across a period, ↓ down a group.
- Electronegativity ↑ across, ↓ down.
- What is the trend in atomic radius across a period?
- Which property measures the tendency of an atom to attract electrons?
- How many valence electrons do elements in Group 17 have?
Periodic Table Elements Preview
Explore the complete periodic table in the full notes!
🧪 No additional files available for this chapter.
Lab Questions
- What are the key chemical principles demonstrated in this chapter?
- How can you apply these concepts to real-world situations?
- What safety precautions should be taken when working with these chemicals?
- Can you predict the outcomes of similar chemical reactions?
CHEMICAL REACTIONS
Types of Chemical Reactions – Visual Guide
Combination
A + B → AB
Example: 2H₂ + O₂ → 2H₂O
Decomposition
AB → A + B
Example: 2HgO → 2Hg + O₂
Displacement
A + BC → AC + B
Example: Zn + CuSO₄ → ZnSO₄ + Cu
Double Displacement
AB + CD → AD + CB
Example: Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl
Combustion
Fuel + O₂ → CO₂ + H₂O
Example: CH₄ + 2O₂ → CO₂ + 2H₂O
Redox
Involves electron transfer
Example: 2Na + Cl₂ → 2NaCl
2H₂ + O₂ → 2H₂O
Hydrogen + Oxygen → Water
🧪 No additional files available for this chapter.
Lab Questions
- What are the key chemical principles demonstrated in this chapter?
- How can you apply these concepts to real-world situations?
- What safety precautions should be taken when working with these chemicals?
- Can you predict the outcomes of similar chemical reactions?
THE MOLE
⚖️ Chapter: The Mole and Molar Mass
1. Introduction to the Mole
In chemistry, atoms and molecules are too small to count individually. The **mole** is the standard unit used to count very large numbers of particles, like atoms, molecules, or ions. It provides a bridge between the atomic scale and the macroscopic, laboratory scale.
2. Avogadro's Number ($N_A$)
The mole is defined as the amount of substance that contains as many elementary entities (atoms, molecules, or other particles) as there are atoms in exactly 12 grams of the isotope carbon-12 ($\text{}^{12}\text{C}$).
- One mole (1 mol) of any substance contains **Avogadro's Number** of particles.
- $$N_A = 6.022 \times 10^{23} \text{ particles/mol}$$
3. Molar Mass (M)
The **Molar Mass (M)** is the mass in grams of one mole of a substance. It is numerically equal to the atomic mass (for elements) or the molecular/formula mass (for compounds), but the units are $\text{g/mol}$.
- For Elements: The molar mass is found directly on the periodic table (e.g., $\text{Na} \approx 22.99 \text{ g/mol}$).
- For Compounds: The molar mass is calculated by summing the atomic masses of all atoms in the chemical formula.
4. Calculation Examples
- Mass of $\text{H} \text{ (2 moles)}: 2 \times 1.01 \text{ g/mol} = 2.02 \text{ g/mol}$
- Mass of $\text{O} \text{ (1 mole)}: 1 \times 16.00 \text{ g/mol} = 16.00 \text{ g/mol}$
- Total Molar Mass ($\text{H}_2\text{O}$): $2.02 + 16.00 = 18.02 \text{ g/mol}$
5. Mole Calculations: The Conversion Factor
The mole acts as the central conversion factor between mass, number of particles, and volume (for gases). The primary relationship is:
$$\text{moles} = \frac{\text{Mass} \text{ (g)}}{\text{Molar Mass} \text{ (g/mol)}}$$
$$\text{moles} = \frac{\text{Number of Particles}}{N_A}$$
[Image of the mole concept map showing conversions]How many moles are in $10.0 \text{ g}$ of $\text{NaCl}$? ($\text{M} = 58.44 \text{ g/mol}$)
$$\text{moles} = \frac{10.0 \text{ g}}{58.44 \text{ g/mol}} \approx 0.171 \text{ mol}$$
6. Review Questions
- Define the mole in your own words.
- What is the value and significance of Avogadro's Number?
- Calculate the molar mass of $\text{CaCO}_3$. (Atomic masses: $\text{Ca}=40.08, \text{C}=12.01, \text{O}=16.00$)
- How many grams of $\text{CO}_2$ are required to have $0.50 \text{ mol}$ of the gas?
🧪 No additional files available for this chapter.
Lab Questions
- What are the key chemical principles demonstrated in this chapter?
- How can you apply these concepts to real-world situations?
- What safety precautions should be taken when working with these chemicals?
- Can you predict the outcomes of similar chemical reactions?
CHEMICAL BONDING
🤝 Chapter: Chemical Bonding
1. Introduction to Bonding
Chemical bonds are the forces that hold atoms together to form compounds. Atoms bond to achieve a more stable electron configuration, usually by satisfying the **octet rule** (having eight valence electrons).
2. Ionic Bonding
**Ionic bonding** involves the complete **transfer** of one or more valence electrons from one atom to another, resulting in the formation of ions.
- **Formation:** Typically occurs between a **metal** (which loses electrons to become a positively charged cation) and a **non-metal** (which gains electrons to become a negatively charged anion).
- **Force:** The resulting ions are held together by a strong electrostatic attraction (opposite charges attract).
- **Example:** Sodium Chloride ($\text{NaCl}$): $\text{Na}$ loses an electron to form $\text{Na}^{+}$, and $\text{Cl}$ gains an electron to form $\text{Cl}^{-}$.
3. Covalent Bonding
**Covalent bonding** involves the **sharing** of valence electrons between atoms.
- **Formation:** Typically occurs between two **non-metals**.
- **Types:**
- **Nonpolar Covalent:** Electrons are shared equally (e.g., $\text{O}_2$).
- **Polar Covalent:** Electrons are shared unequally due to a difference in electronegativity (e.g., $\text{H}_2\text{O}$). This results in partial positive ($\delta+$) and partial negative ($\delta-$) poles.
4. Determining Bond Type
The difference in **electronegativity ($\Delta\text{EN}$)** between two atoms is used to predict the bond type.
| $\Delta\text{EN}$ Range | Bond Type |
|---|---|
| $0.0 - 0.4$ | Nonpolar Covalent |
| $0.5 - 1.7$ | Polar Covalent |
| $> 1.7$ | Ionic |
5. Review Questions
- What is the driving force behind chemical bonding?
- Compare and contrast the electron behavior in ionic vs. covalent bonds.
- If the electronegativity difference is $1.0$, what type of bond is formed?
🧪 No additional files available for this chapter.
Lab Questions
- What are the key chemical principles demonstrated in this chapter?
- How can you apply these concepts to real-world situations?
- What safety precautions should be taken when working with these chemicals?
- Can you predict the outcomes of similar chemical reactions?
Chemical Reactions & Stoichiometry
⚗️ Chapter: Chemical Reactions & Stoichiometry
1. Chemical Equations
A **chemical equation** is a representation of a chemical reaction using chemical formulas. It shows the **reactants** (starting materials) on the left and the **products** (final substances) on the right, separated by an arrow ($\to$).
$$\text{Reactants} \to \text{Products}$$
Example: $$2\text{H}_2 \text{ (g)} + \text{O}_2 \text{ (g)} \to 2\text{H}_2\text{O} \text{ (l)}$$
2. Balancing Equations
Chemical equations must obey the **Law of Conservation of Mass**, meaning atoms are neither created nor destroyed. The number of atoms of each element must be the same on both sides of the equation. This is achieved by adjusting the **coefficients** (numbers placed in front of the formulas).
3. Types of Reactions
| Type | Pattern | Example |
|---|---|---|
| Synthesis (Combination) | $\text{A} + \text{B} \to \text{AB}$ | $2\text{Na} + \text{Cl}_2 \to 2\text{NaCl}$ |
| Decomposition | $\text{AB} \to \text{A} + \text{B}$ | $2\text{H}_2\text{O} \to 2\text{H}_2 + \text{O}_2$ |
| Single Replacement | $\text{A} + \text{BC} \to \text{AC} + \text{B}$ | $\text{Zn} + 2\text{HCl} \to \text{ZnCl}_2 + \text{H}_2$ |
| Double Replacement | $\text{AB} + \text{CD} \to \text{AD} + \text{CB}$ | $\text{AgNO}_3 + \text{NaCl} \to \text{AgCl} + \text{NaNO}_3$ |
4. Stoichiometry: The Mole Bridge
**Stoichiometry** is the area of chemistry that deals with the quantitative relationships between reactants and products in a balanced chemical equation. The **coefficients** in the balanced equation represent the **mole ratio**.
For the reaction $2\text{H}_2 + \text{O}_2 \to 2\text{H}_2\text{O}$, the mole ratio is: $$\frac{2 \text{ mol } \text{H}_2}{1 \text{ mol } \text{O}_2} \quad \text{or} \quad \frac{2 \text{ mol } \text{H}_2\text{O}}{1 \text{ mol } \text{O}_2}$$
5. Review Questions
- State the Law of Conservation of Mass in terms of a chemical equation.
- Balance the following equation: $\text{Fe} + \text{O}_2 \to \text{Fe}_2\text{O}_3$
- In the decomposition of water, what is the mole ratio of hydrogen gas to oxygen gas?
2H₂ + O₂ → 2H₂O
Hydrogen + Oxygen → Water
🧪 No additional files available for this chapter.
Lab Questions
- What are the key chemical principles demonstrated in this chapter?
- How can you apply these concepts to real-world situations?
- What safety precautions should be taken when working with these chemicals?
- Can you predict the outcomes of similar chemical reactions?
Acids and Bases HTML
🧪 Chapter: Acids and Bases
1. Definitions of Acids and Bases
Acids and bases are two fundamental classes of chemical compounds defined by their behavior in water. We primarily use the **Arrhenius** and **Brønsted-Lowry** definitions.
| Type | Arrhenius Definition | Brønsted-Lowry Definition |
|---|---|---|
| Acid | Produces hydrogen ions ($\text{H}^{+}$) in water. | Proton ($\text{H}^{+}$) **donor**. |
| Base | Produces hydroxide ions ($\text{OH}^{-}$) in water. | Proton ($\text{H}^{+}$) **acceptor**. |
*Note: The $\text{H}^{+}$ ion in water exists as the hydronium ion ($\text{H}_3\text{O}^{+}$).*
2. The pH Scale
The **pH scale** is a logarithmic scale used to specify the acidity or basicity of an aqueous solution. It is based on the concentration of the hydronium ion.
- $$\text{pH} = - \log [\text{H}^+]$$
- $$\text{pOH} = - \log [\text{OH}^-]$$
- $$\text{pH} + \text{pOH} = 14$$
3. Interpreting the pH Scale
| pH Value | Classification | $[\text{H}^+]$ vs. $[\text{OH}^-]$ |
|---|---|---|
| $0$ to $< 7$ | Acidic | $[\text{H}^+] > [\text{OH}^-]$ |
| $= 7$ | Neutral | $[\text{H}^+] = [\text{OH}^-]$ |
| $> 7$ to $14$ | Basic (Alkaline) | $[\text{H}^+] < [\text{OH}^-]$ |
4. Neutralization Reactions
A **neutralization reaction** is a type of double replacement reaction where an acid and a base react to form a **salt** and **water**.
$$\text{Acid} + \text{Base} \to \text{Salt} + \text{Water}$$ $$\text{HCl} + \text{NaOH} \to \text{NaCl} + \text{H}_2\text{O}$$
5. Review Questions
- Define an acid and a base according to the Brønsted-Lowry theory.
- A solution has a $[\text{H}^+]$ concentration of $1.0 \times 10^{-4} \text{ M}$. Calculate its pH. Is it acidic or basic?
- Identify the salt formed when $\text{H}_2\text{SO}_4$ reacts with $\text{KOH}$.
🧪 No additional files available for this chapter.
Lab Questions
- What are the key chemical principles demonstrated in this chapter?
- How can you apply these concepts to real-world situations?
- What safety precautions should be taken when working with these chemicals?
- Can you predict the outcomes of similar chemical reactions?
Acids and Bases HTML
🧪 Chapter: Acids and Bases
1. Definitions of Acids and Bases
Acids and bases are two fundamental classes of chemical compounds defined by their behavior in water. We primarily use the **Arrhenius** and **Brønsted-Lowry** definitions.
| Type | Arrhenius Definition | Brønsted-Lowry Definition |
|---|---|---|
| Acid | Produces hydrogen ions ($\text{H}^{+}$) in water. | Proton ($\text{H}^{+}$) **donor**. |
| Base | Produces hydroxide ions ($\text{OH}^{-}$) in water. | Proton ($\text{H}^{+}$) **acceptor**. |
*Note: The $\text{H}^{+}$ ion in water exists as the hydronium ion ($\text{H}_3\text{O}^{+}$).*
2. The pH Scale
The **pH scale** is a logarithmic scale used to specify the acidity or basicity of an aqueous solution. It is based on the concentration of the hydronium ion.
- $$\text{pH} = - \log [\text{H}^+]$$
- $$\text{pOH} = - \log [\text{OH}^-]$$
- $$\text{pH} + \text{pOH} = 14$$
3. Interpreting the pH Scale
| pH Value | Classification | $[\text{H}^+]$ vs. $[\text{OH}^-]$ |
|---|---|---|
| $0$ to $< 7$ | Acidic | $[\text{H}^+] > [\text{OH}^-]$ |
| $= 7$ | Neutral | $[\text{H}^+] = [\text{OH}^-]$ |
| $> 7$ to $14$ | Basic (Alkaline) | $[\text{H}^+] < [\text{OH}^-]$ |
4. Neutralization Reactions
A **neutralization reaction** is a type of double replacement reaction where an acid and a base react to form a **salt** and **water**.
$$\text{Acid} + \text{Base} \to \text{Salt} + \text{Water}$$ $$\text{HCl} + \text{NaOH} \to \text{NaCl} + \text{H}_2\text{O}$$
5. Review Questions
- Define an acid and a base according to the Brønsted-Lowry theory.
- A solution has a $[\text{H}^+]$ concentration of $1.0 \times 10^{-4} \text{ M}$. Calculate its pH. Is it acidic or basic?
- Identify the salt formed when $\text{H}_2\text{SO}_4$ reacts with $\text{KOH}$.
🧪 No additional files available for this chapter.
Lab Questions
- What are the key chemical principles demonstrated in this chapter?
- How can you apply these concepts to real-world situations?
- What safety precautions should be taken when working with these chemicals?
- Can you predict the outcomes of similar chemical reactions?
GAS LAWS
🎈 Chapter: The Gas Laws
1. Properties of Gases
Gases have no definite shape or volume and are highly compressible. Their behavior is described by four interrelated variables: **Volume (V)**, **Pressure (P)**, **Temperature (T)**, and **Moles (n)**.
- Temperature: $$0^\circ \text{C}$ or $273.15 \text{ K}$$
- Pressure: $$1 \text{ atm}$ or $101.3 \text{ kPa}$$
- At STP, $$1 \text{ mole}$ of any gas occupies $22.4 \text{ L}$$ (molar volume).
2. Ideal Gas Law
The **Ideal Gas Law** combines the relationships between P, V, T, and n into a single equation. It describes the behavior of an ideal gas, which is a hypothetical gas whose molecules exhibit no intermolecular forces.
$$\text{PV} = \text{nRT}$$
Where **R** is the **Ideal Gas Constant** ($$0.0821 \frac{\text{L}\cdot\text{atm}}{\text{mol}\cdot\text{K}}$$ or $8.314 \frac{\text{J}}{\text{mol}\cdot\text{K}}$$).
3. The Combined Gas Law (Constant Moles)
The **Combined Gas Law** relates the changes in P, V, and T for a fixed amount of gas (constant moles, $$n$$).
$$\frac{\text{P}_1\text{V}_1}{\text{T}_1} = \frac{\text{P}_2\text{V}_2}{\text{T}_2}$$
4. Individual Gas Laws (Special Cases)
| Law | Relationship | Fixed Variables |
|---|---|---|
| Boyle's Law | $$\text{P} \propto 1/\text{V}$$ (Inverse) | n, T |
| Charles's Law | $$\text{V} \propto \text{T}$$ (Direct) | n, P |
| Gay-Lussac's Law | $$\text{P} \propto \text{T}$$ (Direct) | n, V |
| Avogadro's Law | $$\text{V} \propto \text{n}$$ (Direct) | P, T |
*Crucial Note: All temperatures in gas law calculations must be in **Kelvin (K)**.*
5. Review Questions
- What are the four variables used to describe the state of a gas?
- A balloon holds $5.0 \text{ L}$$ of air at $$25^\circ \text{C}$$. If the temperature is doubled, what happens to the volume (assuming constant pressure)? What law applies?
- Calculate the pressure exerted by $1.5 \text{ mol}$$ of gas in a $$10.0 \text{ L}$$ container at $$300 \text{ K}$$. (Use $$\text{R} = 0.0821 \frac{\text{L}\cdot\text{atm}}{\text{mol}\cdot\text{K}}$$)
🧪 No additional files available for this chapter.
Lab Questions
- What are the key chemical principles demonstrated in this chapter?
- How can you apply these concepts to real-world situations?
- What safety precautions should be taken when working with these chemicals?
- Can you predict the outcomes of similar chemical reactions?
GAS LAWS
🎈 Chapter: The Gas Laws
1. Properties of Gases
Gases have no definite shape or volume and are highly compressible. Their behavior is described by four interrelated variables: **Volume (V)**, **Pressure (P)**, **Temperature (T)**, and **Moles (n)**.
- Temperature: $0^\circ \text{C}$ or $273.15 \text{ K}$
- Pressure: $1 \text{ atm}$ or $101.3 \text{ kPa}$
- At STP, $1 \text{ mole}$ of any gas occupies $22.4 \text{ L}$ (molar volume).
2. Ideal Gas Law
The **Ideal Gas Law** combines the relationships between P, V, T, and n into a single equation. It describes the behavior of an ideal gas, which is a hypothetical gas whose molecules exhibit no intermolecular forces.
$$\text{PV} = \text{nRT}$$
Where **R** is the **Ideal Gas Constant** ($0.0821 \frac{\text{L}\cdot\text{atm}}{\text{mol}\cdot\text{K}}$ or $8.314 \frac{\text{J}}{\text{mol}\cdot\text{K}}$).
3. The Combined Gas Law (Constant Moles)
The **Combined Gas Law** relates the changes in P, V, and T for a fixed amount of gas (constant moles, $n$).
$$\frac{\text{P}_1\text{V}_1}{\text{T}_1} = \frac{\text{P}_2\text{V}_2}{\text{T}_2}$$
4. Individual Gas Laws (Special Cases)
| Law | Relationship | Fixed Variables |
|---|---|---|
| Boyle's Law | $\text{P} \propto 1/\text{V}$ (Inverse) | n, T |
| Charles's Law | $\text{V} \propto \text{T}$ (Direct) | n, P |
| Gay-Lussac's Law | $\text{P} \propto \text{T}$ (Direct) | n, V |
| Avogadro's Law | $\text{V} \propto \text{n}$ (Direct) | P, T |
*Crucial Note: All temperatures in gas law calculations must be in **Kelvin (K)**.*
5. Review Questions
- What are the four variables used to describe the state of a gas?
- A balloon holds $5.0 \text{ L}$ of air at $25^\circ \text{C}$. If the temperature is doubled, what happens to the volume (assuming constant pressure)? What law applies?
- Calculate the pressure exerted by $1.5 \text{ mol}$ of gas in a $10.0 \text{ L}$ container at $300 \text{ K}$. (Use $\text{R} = 0.0821 \frac{\text{L}\cdot\text{atm}}{\text{mol}\cdot\text{K}}$)
🧪 No additional files available for this chapter.
Lab Questions
- What are the key chemical principles demonstrated in this chapter?
- How can you apply these concepts to real-world situations?
- What safety precautions should be taken when working with these chemicals?
- Can you predict the outcomes of similar chemical reactions?