Chapter 1: Matter and Measurements – GOB Chemistry Study Notes

Chapter 1: Matter and Measurements

Chemistry: The Central Science

Chemistry is the study of the nature, properties, and transformations of matter. It connects to many scientific and health-related disciplines, making it a central science in understanding the physical world and biological systems.

• Matter: Anything that has mass and occupies space.

• Applications: Chemistry is foundational to biology, physics, medicine, environmental science, and more.

States of Matter

Matter exists in three primary states: solid, liquid, and gas. The state depends on temperature and energy changes.

• Solid: Definite shape and volume.

• Liquid: Definite volume, variable shape.

• Gas: Variable shape and volume.

• Physical Change: Changes in state are physical changes; the chemical identity remains the same.

Classification of Matter

Matter can be classified based on its composition and uniformity. The main categories are pure substances and mixtures.

• Pure Substance: Has a constant composition; can be an element or a compound.

• Mixture: Contains two or more substances; can be homogeneous (uniform) or heterogeneous (non-uniform).

Classification Examples

• Element: Pure copper wire is an element; it cannot be broken down chemically.

• Compound: Pure table sugar (sucrose) is a compound; it can be broken down chemically.

• Homogeneous Mixture: Salt completely dissolved in water forms a homogeneous mixture.

Physical vs Chemical Changes

Changes in matter are classified as physical or chemical based on whether the substance's identity changes.

• Physical Change: Alters form, size, or state without changing chemical composition (e.g., melting ice, boiling water).

• Chemical Change: Produces new substances with different chemical identities (e.g., burning paper, rusting iron).

States of Matter (Revisited)

Solids, liquids, and gases differ in their properties and behavior.

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes shape of container.

• Gas: No definite shape or volume.

Chemical Elements and Symbols

Elements are represented by unique symbols, usually one or two letters. These symbols are used in chemical formulas and equations.

• Element Symbol: First letter capitalized, second letter lowercase (e.g., C for carbon, Cl for chlorine).

• Common Elements: H, C, N, O, Na, Mg, Al, K, Ca, Fe, Cl, Cu, Zn, Ag, Sn, Au.

Elements of Life and Composition

Elements are distributed differently in the Earth's crust and the human body. Some are essential for life.

Earth's Crust	Human Body
Oxygen 46.1%	Oxygen 61%
Silicon 28.2%	Carbon 23%
Aluminum 8.2%	Hydrogen 10%
Iron 5.6%	Nitrogen 2.6%
Calcium 4.1%	Calcium 1.4%
Sodium 2.1%	Phosphorus 1.1%
Magnesium 2.3%	Sulfur 0.20%
Potassium 2.1%	Potassium 0.20%
Titanium 0.57%	Sodium 0.14%
Hydrogen 0.14%	Chlorine 0.12%

Essential Elements in the Human Body

Some elements are required for life, either as macronutrients (needed in larger amounts) or micronutrients (needed in smaller amounts).

• Macronutrients: C, H, O, N

• Micronutrients: Fe, I, Ca, Na, K, etc.

Chemical Formulas and Subscripts

Chemical formulas use element symbols and subscripts to indicate the types and numbers of atoms in a compound.

• Subscript: Indicates the number of atoms of the element immediately before it.

• Example: H2O has 2 hydrogen atoms and 1 oxygen atom.

Counting Atoms in Formulas

• H2SO4: 2 H, 1 S, 4 O (7 atoms total)

• C6H12O6: 6 C, 12 H, 6 O (24 atoms total)

• Ca3N2: 3 Ca, 2 N (5 atoms total)

Chemical Reactions and Equations

Chemical reactions transform reactants into products, often producing new substances with different properties.

• Reactant: Starting substance

• Product: Substance formed

• Equation: Reactants → Products

Example Reaction

• Decomposition of Water: $2H_2O(l) \rightarrow 2H_2(g) + O_2(g)$

Scientific Notation

Scientific notation is used to express very large or very small numbers efficiently, making calculations easier and clarifying significant figures.

• Format: Coefficient × 10Exponent

• Example: 6.02 × 1023

• Conversion: Move the decimal point to create a coefficient between 1 and 10; exponent shows the number of places moved.

Measurements and Dimensional Analysis

Measurements require both a number and a unit. Dimensional analysis uses conversion factors to change units and solve problems.

• Mass: kg, g, mg, µg

• Length: km, m, cm, mm

• Volume: L, mL, µL, cm3

• Metric Prefixes: kilo (k), milli (m), micro (µ)

Unit Conversion Example

• Convert 2.35 L to mL: $2.35 \times \frac{1000\ mL}{1\ L} = 2350\ mL$

Significant Figures

Significant figures reflect the precision of a measurement. The rules for counting significant figures depend on the presence of zeros and decimal points.

• Nonzero digits and interior zeros: Always count.

• Leading zeros: Do not count.

• Trailing zeros: Count if after a decimal point.

• Exact numbers: Have unlimited significant figures.

Examples

• 94.072 g: 5 significant figures

• 0.0834 cm: 3 significant figures

• 138.200 m: 6 significant figures

• 23,000 kg: Ambiguous; use scientific notation to clarify

Calculations with Significant Figures

The operation determines how the final answer is rounded.

• Multiplication/Division: Round to the same number of significant figures as the factor with the fewest.

• Addition/Subtraction: Round to the least precise decimal place.

Example

• 5.19 cm × 95.01 cm × 0.10 cm = 49 cm3 (2 significant figures)

Temperature, Heat, and Energy

Temperature measures the average kinetic energy of particles; heat is energy transferred due to temperature differences. Three scales are used: Celsius, Fahrenheit, and Kelvin.

• Conversion Equations:

  • $\text{°C} = (\text{°F} - 32) \times \frac{5}{9}$

  • $\text{°F} = (\text{°C} \times \frac{9}{5}) + 32$

  • $K = \text{°C} + 273.15$

• Specific Heat Equation: $q = m \times c \times \Delta T$

Example

• How much heat warms 50.0 g of water from 20.0 °C to 30.0 °C? $q = (50.0\ g)(4.184\ J/(g\cdot\text{°C}))(10.0\text{°C}) = 2092\ J$

Density

Density relates the mass of a substance to the volume it occupies. It predicts whether objects will float or sink in a liquid.

• Equation: $d = \frac{m}{V}$

• Units: g/cm3 (solids), g/mL (liquids), g/L (gases)

• Floating: An object floats if its density is less than the liquid's density.

Example

• Find the density of a liquid with mass 45.0 g and volume 18.0 mL: $d = \frac{45.0\ g}{18.0\ mL} = 2.50\ g/mL$

Scientific Method

The scientific method uses evidence and logical reasoning to explain observations, such as why some canned drinks float.