Introduction to Chemistry: Matter and Composition

• Chemistry studies matter, its transformations, physical, and chemical properties, including changes due to chemical reactions
• The focus is on how matter interacts at the atomic/molecular level

States of Matter

• Matter exists in three phases: solid, liquid, and gas, e.g., H2O(s), H2O(l), and H2O(g), respectively
• Phase transitions happen at specific pressure and temperature values
• Phase transitions are governed by the properties of atoms and molecules

Matter Composition

• Atoms are the fundamental building blocks of matter
• Each element is composed of only one kind of atom
• Compounds consist of two or more different elements

Pure Substances

• Pure substances possess distinct properties and a defined composition
• Pure substances do not vary from sample to sample, for example, salt

Elements

• Elements cannot be broken down into simpler substances, for example, oxygen gas

Compounds

• Compounds are composed of two or more different elements, such as water
• The elemental composition of a pure compound is always the same.
• This is the Law of Constant Composition (or Law of Definite Proportions)

Mixtures

• Mixtures combine two or more substances, each retaining its chemical identity
• Mixtures can be separated from each other

Types of Mixtures

• Homogeneous mixtures are uniform throughout, for example, vanilla ice cream
• Heterogeneous mixtures do not have uniform composition, properties, and appearance throughout, for example, muesli

Physical Properties

• Physical properties are observed without changing the substance
• Examples include boiling point, density, mass, and volume

Chemical Properties

• Chemical properties are observed when a substance is changed into another substance
• Examples include flammability, corrosiveness, and reactivity with acid

Physical Changes

• Physical changes in matter do not alter the composition of a substance
• Changes of state, temperature, and volume are examples of physical changes

Chemical Changes

• Chemical changes involve the formation of new substances
• Examples include combustion, oxidation, and decomposition

Chemical Reactions (Chemical Change)

• In chemical reactions, reacting substances are converted to new substances

Compounds Decomposition

• Compounds can be broken down into more elemental particles via electrolysis, creating hydrogen gas and oxygen gas from water

Separation of Mixtures: Distillation

• Distillation separates homogeneous mixtures based on boiling point differences

Separation of Mixtures: Filtration

• Filtration separates solid substances from liquids and solutions

Separation of Mixtures: Chromatography

• Chromatography separates substances based on varying solubility in a solvent

Scientific Measurement

• Scientific measurement is used to measure quantitative properties of matter

SI Units

• SI (Système International d'Unités) units are used
• Kilogram (kg) measures mass
• Metre (m) measures length
• Second (s) measures time
• Kelvin (K) measures temperature
• Mole (mol) measures amount of substance
• Ampere (A) measures electric current
• Candela (cd) measures luminous intensity

Metric System Prefixes Conversion

• Prefixes convert base metric units into units appropriate for the item to be measured
• Giga (G) is 10^9
• Mega (M) is 10^6
• Kilo (k) is 10^3
• Deci (d) is 10^-1
• Centi (c) is 10^-2
• Milli (m) is 10^-3
• Micro (µ) is 10^-6
• Nano (n) is 10^-9
• Pico (p) is 10^-12
• Femto (f) is 10^-15

SI Units – Length and Mass

• The SI base unit of length is the metre (m)
• Mass (m) measures the amount of material with the SI base unit of kilogram (kg)

Temperature

• The Kelvin is the SI unit of temperature and is based on the properties of gases
• Absolute temperatures do not have negative Kelvin values
• To convert Celsius to Kelvin, use the formula: K = °C + 273.15

Temperature Scales

• Celsius (°C) is based on the freezing (0°C) and boiling (100°C) points of water
• Kelvin (K) is the SI unit, representing the absolute scale and is measured without a degree sign
• Units of Celsius and Kelvin are equal in magnitude vs the Fahrenheit (°F) (the English system)

Temperature Conversion Formulas

• The formula to convert Fahrenheit to Celsius is: °C = (°F – 32) × 5/9
• The formula to convert Celsius to Kelvin is: K = °C + 273.15
• The formula to convert Celsius to Fahrenheit is: °F = (9/5 × °C) + 32

Volume

• Commonly used metric units for volume are the litre (L) and the millilitre (mL)
• One litre is equivalent to a cube that is 1 dm (decimetre) long on each side
• One millilitre is equivalent to a cube that is 1 cm (centimetre) long on each side
• 1 dm³ = (1 dm) x (1 dm) x (1 dm) = 10 cm x 10 cm x 10 cm
• 1000 cm³ = 1 L

Density

• Density is a physical property that indicates the ratio of mass to volume for a substance
• Formula for density: ρ = m/V
• ρ represents density (g/mL)
• m represents mass (g)
• V represents volume (mL or cm³)

Scientific Notation

• Scientific notation is used to express very large or very small numbers
• A sample written number, expressing the width of a human hair (0.000008 m) is 8 x 10-6 m
• A written large number of 4,500,000 is written as 4.5 x 10^6

Components of Scientific Notation

• Numbers in scientific notation include a coefficient and a power of 10
• Writing a number in scientific notation requires the decimal point to be moved after the first digit
• Number of spaces moved is represented by a positive or negative power of ten accordingly

Significant Digits

• The length reported is reported as 2.76 cm
• Digits 2 and 7 are certain/known
• The final digit, 6, is an estimated digit (uncertain)
• All three digits form a significant number, including the estimated digit

Measurement

• Measuring devices vary in their uses and degrees of accuracy

Significant Figures Meaning

• Digits of a measured quantity, including the uncertain one, are called significant figures
• A greater number of significant figures indicates more certainty

Guidelines: Determining Significant Figures

• All nonzero digits are significant, evidenced by e.g. 123.45 contains 5 significant figures
• Zeros between two significant figures are significant, for example, evidenced by 103.405 contains 6 significant figures
• Zeros at the beginning of a number are never significant examples is 00123.45 = 123.45 which contains 5 significant numbers

Rounding Calculations

• Measured numbers must have the same number of significant figures as the answers
• Calculated answers are usually rounded off to conform to the data
• Specific rounding rules are applied to obtain the figures

Rounding Procedure Rules

• Drop the digit if it's less than 5, for example, 8.724 becomes 8.72
• Add 1 to the preceding digit if it's equal to or greater than 5, for example, 8.727 becomes 8.73

Addition and Subtraction Rounding Rule

• Round answers to the least significant decimal place.

Multiplication and Division Rounding Rule

• Round answers to the number of digits that corresponds to the least number of significant figures used

Exact Numbers

• Exact numbers have an infinite number of significant figures
• 3 x 2.5 g = 7.5 g is an example of an exact number that does not limit the number of significant figures in the result.
• For two measured lengths: 6.64 cm and 6.68cm = (6.64 + 6.68) ÷ 2 = 6.66 cm

Accuracy

• Accuracy indicates how close a measurement is to the true value of a quantity

Precision

• Precision indicates the closeness of multiple measurements to each other