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Chapter 1: Matter, Energy, and Measurement

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

    Abbreviation

    Meaning

    Giga

    G

    109

    Mega

    M

    106

    kilo

    k

    103

    centi

    c

    10-2

    milli

    m

    10-3

    micro

    μ

    10-6

    nano

    n

    10-9

Prefix Examples

Example 1: Convert 0.0077 kg to g

Example 2: Convert 8800 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 is 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 for temperature is Kelvin.

  • Absolute zero is equivalent to:

    • 0 K

    • -273.15 C

  • Equation for converting Celsius to Kelvin:

    • K=C+273.15K = C+273.15

  • Equation for coverting Celsius to Fahrenheit:

    • F=95(C+32)F=\dfrac{9}{5}\left( C+32\right)

  • Equation for converting Fahrenheit to Celsius:

    • C=59(F32)C=\dfrac{5}{9}\left( F-32\right)

Density

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

    • Most common units are g/mL or g/cm3

  • Equation for density → density = mass/volume

    • d=mvd=\dfrac{m}{v}

Density Examples

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 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)

Significant Figures (SF) Examples

Example 1: 73.000 has 5 SF

Example 2: 1400 has 2 SF

Example 3: 0.005090 has 4 SF

Example 4: 6.378 × 108 has 4 SF

Example 5: 390200 has 4 SF

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.

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

S

Chapter 1: Matter, Energy, and Measurement

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

    Abbreviation

    Meaning

    Giga

    G

    109

    Mega

    M

    106

    kilo

    k

    103

    centi

    c

    10-2

    milli

    m

    10-3

    micro

    μ

    10-6

    nano

    n

    10-9

Prefix Examples

Example 1: Convert 0.0077 kg to g

Example 2: Convert 8800 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 is 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 for temperature is Kelvin.

  • Absolute zero is equivalent to:

    • 0 K

    • -273.15 C

  • Equation for converting Celsius to Kelvin:

    • K=C+273.15K = C+273.15

  • Equation for coverting Celsius to Fahrenheit:

    • F=95(C+32)F=\dfrac{9}{5}\left( C+32\right)

  • Equation for converting Fahrenheit to Celsius:

    • C=59(F32)C=\dfrac{5}{9}\left( F-32\right)

Density

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

    • Most common units are g/mL or g/cm3

  • Equation for density → density = mass/volume

    • d=mvd=\dfrac{m}{v}

Density Examples

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 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)

Significant Figures (SF) Examples

Example 1: 73.000 has 5 SF

Example 2: 1400 has 2 SF

Example 3: 0.005090 has 4 SF

Example 4: 6.378 × 108 has 4 SF

Example 5: 390200 has 4 SF

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.

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