Chapter 1: Matter and Measurements – GOB Chemistry Study Notes

Matter and Measurements

Chemistry: The Central Science

Chemistry is known as the "central science" because it connects and underpins all other scientific disciplines. It focuses on the study of matter—anything that has mass and occupies space—and the changes it undergoes.

• Matter: Physical material of the universe; has mass and occupies space.

• Scientific Method: Systematic process involving observation, hypothesis, and experimentation to expand knowledge.

• Property: Characteristic useful for identifying substances (e.g., size, color, temperature, composition, reactivity).

• Physical Change: Alters physical form without changing chemical identity (e.g., melting, dissolving).

• Chemical Change: Alters chemical composition, forming new substances (e.g., burning, rusting).

Example: Dissolving sugar in water is a physical change; heating sugar to make caramel is a chemical change.

States of Matter

Matter exists in three primary states, each with distinct properties:

• Solid (s): Definite shape and volume.

• Liquid (l): Definite volume, takes shape of container.

• Gas (g): No definite shape or volume; expands to fill container.

Change of State: Conversion between solid, liquid, and gas (e.g., melting, boiling).

Example: At room temperature, formaldehyde is a gas because its boiling point is below room temperature.

Classification of Matter

All matter can be classified as either pure substances or mixtures:

• Pure Substance: Uniform chemical composition throughout.

• Mixture: Blend of two or more substances, each retaining its identity.

• Homogeneous Mixture: Uniform composition (e.g., salt water).

• Heterogeneous Mixture: Non-uniform composition (e.g., salad).

• Element: Fundamental substance that cannot be broken down chemically.

• Compound: Pure substance that can be broken down into simpler substances by chemical reactions.

Example: Vanilla ice cream is a homogeneous mixture; sugar is a compound.

Chemical Elements and Symbols

There are 118 known elements, 91 of which occur naturally. Chemists use symbols (one or two letters) to represent elements.

• First letter capitalized; second letter (if any) lowercase.

• Some symbols are based on Latin names (e.g., Na for sodium from natrium).

• Chemical Formula: Notation using element symbols and subscripts to show atom ratios in compounds.

Element	Symbol
Hydrogen	H
Oxygen	O
Sodium	Na
Carbon	C
Nitrogen	N
Calcium	Ca
Iron	Fe
Potassium	K
Chlorine	Cl
Phosphorus	P
Magnesium	Mg
Sulfur	S

Example: H2O is the chemical formula for water, indicating two hydrogen atoms and one oxygen atom per molecule.

Chemical Reactions: Examples of Chemical Change

Chemical reactions are processes in which substances (reactants) are transformed into new substances (products).

• Chemical Reaction: Change in identity and composition of substances.

• Reactant: Starting material in a reaction.

• Product: Substance formed as a result of the reaction.

• Reactions are represented with reactants on the left, products on the right, and an arrow indicating the direction of change.

Example: Electrolysis of water produces hydrogen and oxygen gases.

Physical Quantities: Units and Scientific Notation

Physical properties are measured using standardized units. The SI (International System of Units) is the standard in science.

• Mass: Kilogram (kg) in SI; gram (g) in metric.

• Length: Meter (m).

• Volume: Cubic meter (m3) in SI; liter (L) in metric.

• Temperature: Kelvin (K) in SI; Celsius (°C) in metric.

• Time: Second (s).

Prefixes are used to indicate multiples or fractions of units (e.g., milli-, micro-, kilo-).

Prefix	Symbol	Multiplier
kilo-	k	103
centi-	c	10-2
milli-	m	10-3
micro-	μ	10-6
nano-	n	10-9

Scientific Notation: Used to express very large or small numbers as a product of a number (1–10) and a power of ten.

• Example: $0.000000120 = 1.20 \times 10^{-7}$

Measuring Mass, Length, and Volume

Different units are used for measuring mass, length, and volume in chemistry:

• Mass: Measured in grams (g) or kilograms (kg).

• Weight: Force due to gravity; not the same as mass.

• Length: Measured in meters (m), centimeters (cm), or millimeters (mm).

• Volume: Measured in liters (L), milliliters (mL), or cubic centimeters (cm3).

1 L = 1000 mL = 1000 cm3

Measurement and Significant Figures

All measurements have some uncertainty. The number of significant figures reflects the precision of a measurement.

• Significant Figures: All certain digits plus one estimated digit.

• Rules:

  • Zeroes between nonzero digits are significant.

  • Leading zeroes are not significant.

  • Trailing zeroes after a decimal point are significant.

  • Trailing zeroes before an implied decimal may or may not be significant.

Example: 0.0076 mL has two significant figures; 3400.0 has five.

Rounding Off Numbers

When performing calculations, results must be rounded to reflect the correct number of significant figures.

• Multiplication/Division: Result has as many significant figures as the least precise measurement.

• Addition/Subtraction: Result has as many decimal places as the least precise measurement.

• Rounding:

  • If the first dropped digit is 4 or less, drop it.

  • If the first dropped digit is 5 or more, round up.

Problem Solving: Unit Conversions and Estimating Answers

Unit conversions are essential in chemistry. The factor-label (dimensional analysis) method is used to convert between units.

• Conversion Factor: Ratio expressing the relationship between two units (e.g., 1 in = 2.54 cm).

• Set up calculations so that unwanted units cancel, leaving the desired unit.

• Always check answers with a ballpark estimate.

Example: To convert 21.5 inches to centimeters: $21.5 \text{ in} \times \frac{2.54 \text{ cm}}{1 \text{ in}} = 54.6 \text{ cm}$

Temperature, Heat, and Energy

All chemical reactions involve energy changes. Temperature is a measure of heat energy, and different scales are used:

• Kelvin (K): SI unit; absolute zero is 0 K.

• Celsius (°C): 0°C is freezing point of water; 100°C is boiling point.

• Fahrenheit (°F): 32°F is freezing point; 212°F is boiling point.

Temperature Conversions:

• $T(\degree C) = \frac{5}{9} [T(\degree F) - 32]$

• $T(\degree F) = \frac{9}{5} T(\degree C) + 32$

• $T(K) = T(\degree C) + 273.15$

Energy Units: Joule (J) in SI; calorie (cal) in metric. 1 cal = 4.184 J.

Specific Heat: Amount of heat needed to raise 1 g of a substance by 1°C.

• Formula: $q = m \times c \times \Delta T$

• Where $q$ = heat (cal or J), $m$ = mass (g), $c$ = specific heat, $\Delta T$ = temperature change (°C)

Density and Specific Gravity

Density relates mass to volume and is a key property for identifying substances.

• Density: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• Units: g/cm3 for solids, g/mL for liquids.

• Less dense substances float on more dense fluids.

• Specific Gravity: Ratio of the density of a substance to the density of water at the same temperature (unitless).

Example: To find the volume of 25.0 g of isopropyl alcohol with density 0.7855 g/mL:

• $\text{Volume} = \frac{25.0 \text{ g}}{0.7855 \text{ g/mL}} = 31.8 \text{ mL}$

Summary Table: Key Physical Quantities and Units

Quantity	SI Unit	Metric Unit
Mass	kilogram (kg)	gram (g)
Length	meter (m)	centimeter (cm)
Volume	cubic meter (m3)	liter (L), milliliter (mL)
Temperature	kelvin (K)	Celsius (°C)
Time	second (s)	second (s)

Additional info: Where the original text referenced figures or tables not fully visible, standard textbook values and examples have been used to complete the explanations and tables.