CHE103 Exam 1 Review: Chapters 1-5 Study Guide

Chapter 1: Chemistry in Our Lives

Matter and Chemicals

Chemistry studies matter, which is anything that has mass and occupies space. Matter can be composed of chemicals, which are substances with a definite composition.

• Matter: Includes solids, liquids, and gases; can be pure substances or mixtures.

• Chemicals: Substances with a specific chemical formula (e.g., water, H2O).

• Example: Air is a mixture of chemicals; water is a pure chemical.

Scientific Method

The scientific method is a systematic approach to investigation.

• Observation: Gathering information using senses or instruments.

• Hypothesis: A proposed explanation based on observations.

• Experiment: Testing the hypothesis under controlled conditions.

• Conclusion: Analyzing results to accept or reject the hypothesis.

• Example: Testing if salt dissolves faster in warm water than cold water.

Place Values and Calculations

Understanding place values is essential for accurate measurement and calculation.

• Place Values: Hundreds, tens, ones, tenths, hundredths, etc.

• Percentage Calculation: $\text{Percentage} = \frac{\text{portion}}{\text{total}} \times 100\%$

• Math Skills: Addition, subtraction, multiplication, division.

• Scientific Notation: Expressing numbers as $a \times 10^n$ (e.g., $3.2 \times 10^4$).

Graph Interpretation

Graphs visually represent data and trends.

• Key Points: Identify axes, units, and trends.

• Example: Temperature vs. time graph for heating water.

Chapter 2: Chemistry and Measurement

Exact vs. Measured Numbers

Numbers in chemistry are either exact (defined values) or measured (obtained by measurement).

• Exact Numbers: Counting numbers or defined conversions (e.g., 1 dozen = 12).

• Measured Numbers: Obtained using instruments; have uncertainty.

Significant Figures (SFs)

Significant figures reflect the precision of a measurement.

• Reporting SFs: Include all certain digits plus one estimated digit.

• Rules:

  • Addition/Subtraction: Keep the fewest decimal places.

  • Multiplication/Division: Keep the fewest SFs.

• Example: $2.45 \times 3.1 = 7.6$ (2 SFs).

Unit Conversions

Converting between units is fundamental in chemistry.

• Metric Prefixes: kilo (k), deci (d), centi (c), milli (m), micro (μ).

• Common Conversions:

  • 1 kg = 1000 g

  • 1 g = 1000 mg

  • 1 L = 1000 mL

  • 1 m = 100 cm

• Conversion Factors: Used to convert units in calculations.

• Example: $\text{cm} \rightarrow \text{mm}$ using $1 \text{cm} = 10 \text{mm}$.

Density

Density is the mass per unit volume of a substance.

• Formula: $\text{Density} = \frac{\text{Mass (g)}}{\text{Volume (mL)}}$

• Example: If mass = 10 g, volume = 2 mL, density = 5 g/mL.

Chapter 3: Matter and Energy

Classification of Matter

Matter can be classified as pure substances or mixtures.

• Pure Substance: Element (e.g., gold) or compound (e.g., water).

• Mixture: Homogeneous (uniform, e.g., salt water) or heterogeneous (non-uniform, e.g., sand and water).

States of Matter

There are three main states of matter, each with distinct properties.

• Solid: Definite shape and volume.

• Liquid: Definite volume, no definite shape.

• Gas: No definite shape or volume.

Chemical vs. Physical Changes and Properties

• Physical Change: Change in state or appearance (e.g., melting ice).

• Chemical Change: Formation of new substances (e.g., rusting iron).

• Physical Property: Observable without changing composition (e.g., boiling point).

• Chemical Property: Describes ability to undergo chemical change (e.g., flammability).

Temperature Conversions

Temperature can be measured in Celsius (°C), Fahrenheit (°F), or Kelvin (K).

• Formulas:

  • $T_F = 1.8 T_C + 32$

  • $T_C = \frac{T_F - 32}{1.8}$

  • $T_K = T_C + 273$

  • $T_C = T_K - 273$

Energy and Specific Heat

Energy is the capacity to do work; it can be potential or kinetic.

• Potential Energy: Stored energy (e.g., chemical bonds).

• Kinetic Energy: Energy of motion.

• Conversions: $1 \text{cal} = 4.184 \text{J}$, $1 \text{kJ} = 1000 \text{J}$, $1 \text{kcal} = 1000 \text{cal}$

• Specific Heat Formula: $\text{SH} = \frac{\text{heat (J)}}{\text{mass (g)} \times \Delta T (\degree C)}$

• Example: Calculate heat required to raise temperature of 10 g water by 5°C.

Energy in Food

• Energy Value: $\text{Energy Value (kJ/g)} = \frac{\text{Energy (kJ)}}{\text{mass (g)}}$

• Example: Calculate total energy from fat, protein, and carbohydrates.

Changes of State

• Melting: Solid to liquid

• Freezing: Liquid to solid

• Boiling: Liquid to gas

• Condensation: Gas to liquid

Heating and Cooling Curves

Heating and cooling curves show temperature changes during phase transitions.

• Melting Point: Temperature at which solid becomes liquid.

• Boiling Point: Temperature at which liquid becomes gas.

Chapter 4: Atoms and Elements

Element Names and Symbols

Elements are represented by unique symbols; knowledge of common elements is essential.

• Periods 1-4: Hydrogen (H) to Krypton (Kr)

• Other Common Elements: Gold (Au), Silver (Ag), Mercury (Hg), Lead (Pb), Tin (Sn), Platinum (Pt)

Periodic Table Structure

• Groups: Vertical columns (e.g., Group 1A, 2A, 7A, 8A)

• Periods: Horizontal rows

• Classification: Metals, nonmetals, metalloids

Atomic Structure and Symbols

• Proton: Positive charge, in nucleus

• Neutron: Neutral, in nucleus

• Electron: Negative charge, outside nucleus

• Atomic Number (Z): Number of protons

• Mass Number (A): Protons + neutrons

• Atomic Symbol: $\text{Element}^{A}_{Z}$

Isotopes

• Definition: Atoms of the same element with different numbers of neutrons.

• Example: $^{12}\text{C}$ and $^{13}\text{C}$

Calculating Subatomic Particles

• Protons: Equal to atomic number

• Neutrons: Mass number - atomic number

• Electrons: Equal to protons in neutral atom

Electron Arrangement and Valence Electrons

• First 20 Elements: Electron arrangement follows energy levels (shells)

• Valence Electrons: Electrons in outermost shell; determine chemical reactivity

• Lewis Symbol: Element symbol with dots representing valence electrons

Periodic Trends

• Atomic Size: Increases down a group, decreases across a period

• Metallic Character: Increases down a group, decreases across a period

• Ionization Energy: Energy required to remove an electron; increases across a period

Chapter 5: Nuclear Chemistry

Types of Radiation

Nuclear chemistry involves radioactive decay and types of radiation.

• Alpha Particle: $^4_2\text{He}$; mass number 4, atomic number 2

• Beta Particle: $^0_{-1}\text{e}$; mass number 0, atomic number -1

• Positron: $^0_{+1}\text{e}$; mass number 0, atomic number +1

• Gamma Ray: $^0_0\gamma$; mass number 0, atomic number 0

Biological Effects and Protection

• Effects: Radiation can damage cells and DNA.

• Protection: Use shielding (lead, concrete), limit exposure time, increase distance.

Radioactive Decay Equations

• Balanced Nuclear Equation: Shows mass and atomic numbers before and after decay.

• Example: $^{14}_6\text{C} \rightarrow ^{14}_7\text{N} + ^0_{-1}\text{e}$ (beta decay)

Half-Life Calculations

• Half-Life: Time required for half the radioactive atoms to decay.

• Amount Remaining: $\text{Remaining} = \text{Initial} \times (\frac{1}{2})^{n}$, where $n$ = number of half-lives

• Number of Half-Lives: $n = \frac{\log(\text{Final}/\text{Initial})}{\log(1/2)}$

Summary of Key Equations and Conversion Factors

Conversion	Equality
Length	1 ft = 12 in.; 1 in. = 2.54 cm; 1 m = 39.4 in.; 1 km = 1000 m; 1 m = 100 cm; 1 cm = 10 mm; 1 m = 1000 mm
Mass	1 kg = 2.20 lb; 1 kg = 1000 g; 1 g = 1000 mg; 1 mg = 1000 mcg
Volume	1 L = 1000 mL; 1 mL = 1 cc = 1 cm3
Time	1 hour = 60 min; 1 min = 60 s
Energy	1 cal = 4.184 J; 1 kJ = 1000 J; 1 kcal = 1000 cal
Temperature	$T_F = 1.8 T_C + 32$; $T_C = \frac{T_F - 32}{1.8}$; $T_K = T_C + 273$; $T_C = T_K - 273$
Density	$\text{Density} = \frac{\text{Mass (g)}}{\text{Volume (mL)}}$
Specific Heat	$\text{SH} = \frac{\text{heat (J)}}{\text{mass (g)} \times \Delta T (\degree C)}$
Percentage	$\text{Percentage} = \frac{\text{portion}}{\text{total}} \times 100\%$
Energy Value	$\text{Energy Value (kJ/g)} = \frac{\text{Energy (kJ)}}{\text{mass (g)}}$

Additional info: This study guide covers foundational concepts in chemistry, measurement, matter, atomic structure, and nuclear chemistry, as outlined for CHE103 Exam 1. All equations and conversion factors are provided for exam preparation.