1.1 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.

• 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.

1.2 Classifications of Matter

States of Matter

• Three states of matter:

  • Solid (s)

    • Fixed volume and shape.

    • Molecules packed tightly together.

    • Fixed shape and volume.

  • Liquid (l)

    • Fixed volume and shape fits to container.

    • Closely packed molecules.

    • Molecules move fast.

  • 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.

  • 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.

  • Ex: Water and table salt.

• Elements

  • Can’t be decomposed into simpler substances.

  • Composed of only one kind of atom.

• Compounds

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

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

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

• 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.

• Homogeneous mixtures have uniformed compositions.

  • Aka solutions

1.3 Properties of Matter

• 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

• 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.

  • Ex: Water

  • Changes of state are physical changes.

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

  • Ex: Burning, rusting

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.

1.4 The Nature of Energy

• Energy - the ability to do work or transfer heat.

• Work - the energy transferred when a force is exerted on an object causes a displacement of that object.

• Heat - the energy used to cause the temperature of an object to increase.

• Force - any push or pull exerted on an object.

• Equation for work:

  • $w=F\times d$

  • w = work

  • F = force

  • d = distance

Kinetic Energy and Potential Energy

• Kinetic energy - energy of motion.

  • Equation for kinetic energy:

    • $KE=\dfrac{1}{2}mv^{2}$

    • KE = kinetic energy

    • m = mass

    • v = velocity

  • Molecules and atoms are in motion.

• Potential energy - stored energy.

  • Equation for potential energy:

    • $PE=mgh$

    • PE = potential energy

    • m = mass

    • g = gravity

    • h = height

• Electrostatic potential energy - arises from the interactions between charged particles.

  • Opposite charges attract, same charges repel.

1.5 Units of Measurement

• Quantitative - numerical 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	mol
  Electric current	Ampere	A or amp
  Luminous intensity	Candela	cd

The Scientific Method

• Hypothesis - a prediction of what the results of an experiment/research may be based on observations.

• Theory - a model that has predictive powers and accounts for all observations.

• Steps of the scientific method:

  • Collect information

  • Formulate a hypothesis

  • Test the hypothesis

  • Formulate a theory

  • Repeatedly test theory

Length and Mass

• SI unit for length is the meter.

• SI unit for mass is the kilogram

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 for temperature is Kelvin.

  • Absolute zero

    • 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:

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

Derived SI Units

• A derived unit is obtained by multiplication or division of one or more base units.

Volume

• The derived SI unit for volume is m³ (cubic meter).

• Liters are also used.

Density

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

• Equation for density:

  • density = mass/volume

  • $d=\dfrac{m}{v}$

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

1.6 Uncertainty in Measurement

• Exact numbers - exact values

  • Defined values

  • Example: 12 eggs in a dozen

• Inexact numbers - values of some uncertainty

  • Numbers from measurements.

  • Uncertainties always exist in measured quantities.

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

Precision and Accuracy

• Precision - measure of how closely individual measurements agree with one another.

• Accuracy - how closely individual measurements agree with the correct/”true” value.

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 places

      • 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)

1.7 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.

Chapter Equations