The Study of Chemistry

The Atomic and Molecular Perspective of Chemistry

• Chemistry: the study of the properties and behavior of matter.

• Matter: the physical material of the universe; anything that has mass and takes up space.

• A property is any characteristic that allows us to recognize a particular type of matter and to distinguish it from other types.

• Elements: combine together to create matter.

• Atoms: the tiniest particles that are the building blocks of matter and can not be divided further.

• Molecules: two or more atoms.

  • Different molecules can be made from the same elements.

Why Study Chemistry?

• Helps improve pharmaceuticals, fertilizers and pesticides, plastics, solar panels, light emitting diodes, and building materials.

• Identify harmful chemicals.

Chemists

• They do three things:

  • Make new types of matter, materials, substances, or combinations of substances with desired properties.

  • Measure the properties of matter.

  • Develop models that explain and/or predict the properties of matter.

States of Matter

Three states of matter

• Solid (s)

  • Fixed volume and shape.

  • Molecules packed tightly together.

  • Example: Ice

• Liquid (l)

  • Fixed volume and shape fits to container.

  • Closely packed molecules.

  • Example: Liquid water

• Gas (g)

  • No fixed volume or shape, fits container.

  • Molecules are far apart.

  • Molecules can move fast and bounce off container walls.

  • Open space = less interaction between molecules.

  • Smaller container = molecules hit each other.

  • Collisions do not affect shape or volume.

  • Example: Water vapor

• Aqueous (aq)

  • Solid dissolved in liquid.

• States of matter can change through temperature or pressure.

Pure Substances

• Pure substance: matter that has distinct properties and a composition that does not vary from sample to sample.

  • Example: Water and table salt.

• The two type of substances are elements and compounds.

  • Elements

    • Can’t be decomposed into simpler substances.

    • Composed of only one kind of atom.

    • Example: Ar, Be, Xe, C

  • Compounds

    • Can be decomposed because it is made up of two or more elements, so the are two or more types of atoms.

    • Example: CO₂, H₂O, C₂H₄

Elements

• 118 named elements.

• Element symbols have one or two letters. First letter capitalized, second lowercase.

• Elements are found on the periodic table.

• Columns of periodic table have elements with similar properties.

Compounds

• Elements can form compounds.

• Law of constant composition: states that the elemental composition of a compound is always the same.

Mixtures

• Mixtures: combinations of two or more substances in which each substance retains its chemical identity.

  • Two different types: homogeneous and heterogenous

• Matter is made up of mixtures of different substances.

• Mixtures can have various compositions.

• Components of a mixture are substances making up a mixture.

• Heterogeneous mixtures vary in composition.

  • Do not mix evenly

  • Examples: Salad, sand in water (does not dissolve), soil

• Homogeneous mixtures have uniformed compositions (evenly mixed).

  • Evenly mix

  • Aka solutions

  • Examples: Air, sugar is water (dissolves), steel

Two Types of Properties

• Physical properties can be observed without changing the identity and composition of the substance.

  • Color, odor, density, melting point, boiling point, hardness

• Chemical properties describe the way a substance may change, or react, to form other substances.

  • Flammability

• Intensive properties do not depend on the amount of sample being examined and are particularly useful in chemistry.

  • Temperature, melting point

  • These are important for identifying a substance.

• Extensive properties depend on the amount of sample. Relates to the amount of substance present.

  • Mass, volume

Physical and Chemical Changes

• Physical changes: substance changes physical appearance but not composition.

  • Example: Cutting a carrot, water changing phases

  • Changes of state are physical changes.

• Chemical change (aka chemical reaction): a substance transforms into a new substance.

  • Example: Burning, rusting, mixing vinegar and baking soda

Separation of Mixtures

• Filtration: separates solids from liquids or gases using a filter.

• Distillation: a separation process that depends on the different abilities of substances to form gases.

• Chromatography: separates substances on the basis of differences in the ability of substances to adhere to the solid surface, in this case, dyes to paper.

Numbers and Chemistry

• Quantitative: numerical measurements

Units of Measurements

SI Units

• SI units: preferred metric units for science.

  Physical Quantity	Unit Name	Abbrev.
  Length	Meter	m
  Mass	Kilograms	kg
  Temperature	Kelvin	K
  Time	Second	s or sec
  Amount of substance	Mole	mole (mol)
  Electric current	Ampere	A or amp
  Luminous intensity	Candela	cd

Metric System

• The base units used in the metric system:

  Physical Quantity	Unit Name	Abbrev.
  Mass	gram	g
  Length	meter	m
  Time	second	s or sec
  Temperature	Degrees Celsius	°C
  Temperature	Kelvin	K
  Amount of substance	mole	mol
  Volume	cubic meter	cm3
  Volume	liter	l

Prefixes for Measurements

• Different prefixes are used to different values

  Prefix	Abbrev.	Scientific Notation	Value
  Giga	G	1 x 109	1,000,000,000
  Mega	M	1 x 106	1,000,000
  kilo	k	1 x 103	1,000
  centi	c	1 x 10-2	0.01
  milli	m	1 x 10-3	0.001
  micro	μ	1 x 10-6	0.000006
  nano	n	1 x 10-9	0.000000001

Prefix Examples

Canceled units are shown in pink

Example 1: Convert 0.0077 kg to g

Example 2: Convert 8,800 mL to L

Example 3: Convert 0.00450 cm to nm

Mass, Length, and Volume

• Mass: a measure of the amount of material in an object.

  • SI unit: kilogram

• Length: a measure of distance.

  • SI unit: meter

• Volume: a derived unit from length.

  • Equation: L x W x H

  • Common units are the liter and milliliter

  • Example: cm x cm x cm = cm³

Temperature

• Temperature: a measure of the hotness or coldness of an object, is a physical property that determines the direction of heat flow.

• SI unit: Kelvin

• Absolute zero is equivalent to:

  • 0 K = -273.15 C

• Equation for converting Celsius to Kelvin

  • $K = C+273.15$

• Equation for coverting Celsius to Fahrenheit

  • $F=\dfrac{9}{5}\left( C+32\right)$

• Equation for converting Fahrenheit to Celsius

  • $C=\dfrac{5}{9}\left( F-32\right)$

Density

• Density: the amount of mass in a unit volume of a substance.

  • Most common SI units: g/mL or g/cm³

• Equation for density → density = mass/volume

  • $d=\dfrac{m}{v}$

Density Examples

Canceled units are shown in pink

Example 1: Find the density of an object using the given information.

Example 2: Find the volume of an object using the given information.

Example 3: Find the mass of an object using the given information

Numbers in Science

• Exact numbers: exact values

  • Defined values

  • Examples: 12 eggs in a dozen, 12 inches in 1 foot

• Inexact numbers: values of some uncertainty

  • Numbers from measurements.

  • Uncertainties always exist in measured quantities.

  • May be inexact from errors (equipment or human errors).

  • Examples: blood pressure, weight, height

Accuracy vs Precision

• Precision: a measure how closely individual measurements agree with one another.

• Accuracy: how closely individual measurements agree with the correct or “true” value.

• Experimentally, we take several measurements and determine a standard deviation.

  

Significant Figures

• Significant figures: all digits of a measured quantity.

• The greater amount of significant figures, the more precise the measurement is.

• What numbers are significant:

  • All non-zeros

  • Zeros between non-zeros

  • Zeros at the end if theres a decimal point

• Zeros at the beginning of a number are NEVER SIGNIFICANT.

Adding and subtracting significant figures

• The answer has the same number of decimal places as the measurement with the fewest decimal places.

  • 20.42 + 1.322 + 83.1 = 104.842

    • 20.42 = two decimal places

    • 1.322 = three decimal places

    • 83.1 = one decimal place!!

    • Answer: 104.8 (one decimal place)

Multiplying and dividing significant figures

• The answer has the same number of significant figures as the measurement with the fewest significant figures.

  • (6.221)(5.2) = 32.3492

  • 6.221 = four significant figure

  • 5.2 = 2 significant figures!!

  • Answer: 32 (2 significant figures)

Significant Figures (SF) Examples

Example 1: 73.000 has 5 SF

Example 2: 1400 has 2 SF

Example 3: 0.005090 has 4 SF (the zeros before 5 do not count)

Example 4: 6.378 × 10⁸ has 4 SF

Example 5: 390200 has 4 SF (the zeros after 2 do not count, there is no decimal point)

Dimensional Analysis

• In dimensional analysis, units are multiplied together or divided into each other along with the numerical values.

• Equivalent units cancel out.

Conversion Factors

• Conversion factor: a fraction whose numerator and denominator are the same quantity expressed in different units.

  • Examples:

    • 1 foot/12 inches = 12 inches/1 foot

• Denominator is used to cancel units.

• Given unit x Desired unit/Given unit

  • Given unit cancels

• Two or more conversion factors:

  • First conversion cancel given unit.

  • Following conversions cancels another unit and gives desired.

• Know that 1 mL = 1 cm³

Units of Energy

• The SI unit for energy is the joule (J).

  • A larger SI unit used is the kilojoule (kJ).

• A calorie (cal) is a non-SI unit that is the amount of energy required to raise the temperature of 1 g of water by one degree Celsius.

• 1 cal = 4.184 J

• 1 Cal = 1000 cal = 1 kcal