Chapter 1 typed.pptx

General Chemistry 1:
Introduction to
Chemistry

Why DoWe Study Chemistry?
• To be better informed
• To be a knowledgeable consumer
• To make better decisions for yourself and society
• To learn problem-solving skills (critical thinking)
• To enhance analytical thinking

Chemistry as the “Central Science”
• Chemistry = the study of matter and the transformation it undergoes
• EVERYTHING is around us is matter (ie a CHEMICAL)
• Table salt = sodium chloride, NaCl
• Table sugar = sucrose, C12H22O11
• Clothes:Wool?Cotton? Polyester?
• Body: lipids, Proteins, Carbohydrates, DNA/RNA
• You name it– it’s a chemical!

The Study of Chemistry
• Chemistry is everywhere!
• Matter is everywhere!
• Thus, chemistry matters!
• Chemistry involves the study of matter – its properties and behavior.
• Macroscopic observations are rooted in microscopic structure.

Chemistry as the “Central Science”

Learning Goals in this Class
• Understand basic principles of general chemistry.
• The main goal of this course is of course to learn about chemistry.
• Some of the most important topics we will cover include dimensional analysis (unit
conversions), nomenclature (naming of chemicals), chemical reactions (types and why
chemicals react), stoichiometry (calculation of amounts of products and reactants in a
chemical reaction), thermodynamics (heat associated with chemical processes),
chemical bonding, chemical structure (how atoms are attached to other atoms), and the
laws that govern gases.

Learning Goals in this Class
• Be able to explain everyday chemical phenomena and apply your knowledge to
solving real-world problems.
• Why is the sky blue?Why do dome reactions feel hot while others feel cold?
• Learning chemistry and doing problems within chemistry improves your critical thinking
skills, which will help you later in life and in other courses.
• Be a better consumer.
• Be able to read and understand product labels, know if a product work as advertised or is
it a scam, understand if a product may be dangerous or dangerous if combined with
another product.
• Be able to explain why chemistry is important and how it applies to everyday life.
• You should be able to answer such questions as why certain chemicals should never be
mixed, what chemical reaction allows bread to rise, how does Drano work, and more …

Chapter 1: Matter,
Energy, and
Measurement

Classifications of Matter
• Chemistry is the study of matter and the transformations matter undergoes
• Matter is defined as all of the physical material in the universe, anything that has mass
and occupies space (volume).
• Matter can be classified in two ways: by physical state or composition.
• 3 states of matter: gas, liquid, and solid.
• Composition: what matter is made of, whether it is a pure substance (only on type
of material) or a mixture (two or more pure substances mixed together).

Physical State: States of Matter
• Gas: a state of matter in which the particles that make up the gas are very far apart,
leading to a substance that has a relatively low density and viscosity.The particles
within a gas have a very large amount of kinetic energy (the energy of motion), and
move very quickly. Gases assume the shape of their container (indefinite shape), have
a indefinite volume, and can expand and compress easily.
• Liquid: a state of matter in which the particles that make up the liquid move freely
among themselves, but are much closer together than the particles of a gas.The
particles within a liquid have a large amount of kinetic energy, and move fairly
quickly. Liquids assume the shape of their container (indefinite shape), but have a
definite volume; they do not expand or compress readily.
• Solid: a state of matter in which the particles that make up the solid do not move very
much (low kinetic energy and rigid structure). Solids are generally the most dense of
the three states of matter. Solids have a definite shape and volume; they do not
expand or compress readily.

Composition: Elements, Compounds, and
Mixtures
• All matter is composed of the some combination of the same 118 elements (on the
periodic table).
• Pure substances are either elements or compounds:
• Element is a substance that can’t be broken down into a simpler substance via chemical
means.
• Elements are composed of atoms.The element carbon is composed of all the atoms of carbons
within the universe.
• Compounds are composed of two or more different atoms (or elements) chemically
bonded to each other.
• Mixture: a combination of two or more different pure substances

Pure Substances
• The two types of pure substances
can be further broken down.
• Elements can be broken down into
single atoms (monatomic), or
molecules (diatomic or polyatomic
where all the atoms are exactly
the same).
• Compounds can be broken down
into molecules or ions (ionic
compounds). Molecules are
compounds that are made of
nonmetals (covalent compounds)
while ionic compounds are made
of a metal with a nonmetal.

Elements and the PeriodicTable
• There are 118 elements on the periodic table.
• The elemental form of an element is not necessarily how the element is shown on
the periodic table.
• The elemental form is the natural form the element takes when it is not combined
with any other element.
• If something is said to be in its elemental form, it can be either monoatomic,
diatomic, or polyatomic.
• Monoatomic elements = Ne, He, C
• Diatomic molecules = H2 , N2 , O2 , F2 , Cl2 , Br2 all elemental forms
• Polyatomic molecules = S8 , P4

Elemental Names and Symbols
• Each element has its own unique name and elemental (or chemical) symbol.
• Elements within the periodic table are arranged roughly by weight (actually by the
number of protons)
• The elemental symbol is a 1 – 2 letters long, where the first letter is capitalized,
and the second (if applicable) is lower case.

Compound Properties
• Compound properties are very different from their composite elements.
• A specific compound (such as water) will always have the same elemental
composition = Law of Constant Composition or Law of Definite Proportions
• This scientific law was first stated by the French chemist Joseph Louis Proust.

Mixtures
• Two or more pure substances mixed together in some ratio.
• Pure substances have a fixed composition, mixtures do not.
• The components of the mixture (the composite pure substances), are not altered
in any way (no chemical reaction occurs between the pure substances).
• Any mixture can be broken down (separated) into the composite pure substances
via physical means (in theory).

Types of Mixtures: Heterogeneous vs
Homogeneous
• The things that make up a mixture are called the
components.The components are the pure
substances.
• Mixtures can be defined as either heterogeneous or
homogeneous.
• Homogeneous mixtures are ones where you cannot
see the different components of the mixture with the
unaided eye; the mixture appears homogenous.
• Ie: sweet tea, salt water, air
• Heterogeneous mixtures are ones where you can see
the different components that make up the mixture.
• Ie: granite, blood, sea water

Solutions
• Type of homogeneous mixture.
• One, or more components are completely dissolve
(dispersed uniformly) within the other component.
• The component that is in greater quantities is
referred to as the solvent, while the component(s)
in lesser quantities is (are) the solute(s).
• Components of a solution are very small in size.

Properties of Matter
• Matter = physical substance of the universe.
• Each different substance has its own unique properties.
• Property = a characteristic of a substance that allows you to distinguish it from
other substances.
• Properties can be considered either physical or chemical, they can also be
considered either intrinsic or extrinsic.

Physical and Chemical Properties
• Physical properties = observable properties.
• Properties we can observe without changing the chemical identity of the substance.
• Examples: color, odor, density, melting point, boiling point.
• Chemical properties = a property which is only observable during a change in the
chemical identity of the substance.
• Observed during a chemical reaction.
• Examples: flammability, reactivity

Intensive and Extensive Properties
• Intensive properties = inherent properties.
• Properties which exist for a material regardless of amount of substance.
• Used to identify a substance.
• Examples: melting point, boiling point, density, temperature
• Extensive properties = properties which are dependent upon amount of
substance.
• Examples: mass, volume, length

Physical Changes and Changes of State
• A physical change is a change in the appearance of
a substance without changing the chemical identity
of the substance.
• A physical change that is often mistaken for a
chemical change is a change in state.
• Going from ice to liquid water to water vapor, the
chemical identity does not change (always H2O), just
the physical form.
Ice water water vapor
H2O(s) H2O(l) H2O(g)
melting
freezing
evaporation
condensation
melting
freezing
evaporation
condensation

Chemical Changes and Chemical
Reactions
• Chemical change = a change in the chemical identity of the
substance, also called a chemical reaction.
• One substance is transformed into another substance.
• Example: burning methane
CH4 (g) + 2 O2 (g)  2 H2O(l) + CO2 (g)
• Example: iron rusting
4 Fe(s) + 3 O2 (g)  2 Fe2O3 (s)

Energy
• Energy = the capacity to do work or transfer heat.
• Energy is either potential (stored) or kinetic (movement).
• Work = measure of energy transfer that occurs when an object is moved over a
distance by an external force. w = F x d
• Heat = a form of energy associated with the movement of atoms and molecules in
any material.The higher the temperature of a material, the faster the atoms are
moving, and hence the greater the amount of energy present as heat. Heat energy
flows from a warmer substance or object to a colder one.
• Temperature is a measure of the average kinetic energy of a system.
• It is a physical property which tells us the direction of heat flow.
• The rate of heat flow depends on the difference in temperature (greater temperature
difference, greater heat flow).

Potential vs Kinetic Energy
• Potential Energy = stored energy
• Electrostatic potential energy = interaction between charged particles.
• Chemical potential energy = energy stored in arrangement of atoms within a molecule.
• Kinetic energy = energy of motion
• Ek = ½ mv2

Units of Measurement
• Definite magnitude of a quantity.
• Means of expressing a physical quantity in any property of matter which is
quantitative (associated with numbers)
• Measurements WILL ALWAYS HAVE units!!!!

Tools used to Gather Measurements
Thermometer:
Temperature (K)
Ruler:
Length (m)
Stopwatch:
Time (s)
Balance:
Mass (g)
Graduated Cylinder:
Volume (L)

The Metric System
• In science (and most of the world), the units of measurement are in the metric
system (meter, liter, gram).
• The US uses the English system (feet, inches, yards).

SI Units and Derived SI Units
• SI Units are a set of units in the metric system that were decided upon by the
French Systeme International d’Unites.
• Base unit from which every other unit is derived (derived SI units)
• Derived SI units are obtained by multiplying/ dividing one or more of the SI units
together.

Macroscopic, Microscopic,Atomic
• Units can be described as either macroscopic, microscopic, or atomic.
• Macroscopic = can be see with the naked eye (0.1 mm or larger)
• Microscopic = seen using a light microscope ( 0.2 μm – 0.1 mm)
• Atomic = cant be seen using a light microscope (some in this region can be seen
with an electron microscope).
• Atoms, molecules, and very small particles

Prefixes
• Measurements can be expressed as
fractions or multiples of a base unit by
using a set of simple prefixes.
• Prefixes precede a base unit to indicate
the fraction or multiple.

Length and Mass
• Length SI unit = meter
• A measure of how long a substance is
• Mass SI unit = kg
• A measure of the amount of material of a substance.
• Mass is independent of gravity (unlike weight)

Temperature
• Temperature = the hotness or coldness of an object (something cold lacks heat.
• SI unit is Kelvin
• Common unit in science in degrees Celsius.
• 0°C = freezing point of pure water, 100°C boiling point of pure water.

Temperature Scales
• (C * 9/5) + 32
• (F - 32) * 5/9
• C + 273.15
• K - 273.15

Volume
• Derived measurement from length
• Derived SI unit = m3
• Some important conversions:
• 1 dm3 = 1 L
• 1 cm3 = 1 mL

Density
• The amount of mass per volume.
• Density doesn’t usually use a derived SI unit.
• Solids are generally expressed in units of g/cm3 while liquids
are in g/mL.
• More dense = more mass in a given volume.

Precision and Accuracy
• Good measurements are ones who are both precise and accurate.
• Precision: how closely individual measurements agree with one another.
• Accuracy: How close an individual measurement agrees with the correct
or “true” value.
ResultA Result B Result C
Accurate? No Yes No
Precise? Yes Yes No

Trials and Standard Deviation
• To ensure that measurements are as precise and accurate as possible, the
measurements are taken several times, called trials.
• The values that are obtained from the various trials are averaged in order to obtain a
good measurement.
• The precision of the measurement is then expressed in terms of a standard deviation.
• Standard deviation: how much an individual measurement differs or deviates from the
average.
• Even if measurements are precise, they might not necessarily be accurate. If the
tool used to obtain the measurement is incorrectly calibrated, the measurements
will be precise but inaccurate.

Exact and Inexact Numbers
• When recording a measurement, there are two types of numbers, exact and
inexact.
• Exact numbers: those whose values are know exactly. Generally determined by counting
objects.
• 9 people in this class
• 12 eggs in 1 dozen
• 1000 g in 1 kilogram the 1’s in all conversions are also considered exact
• 2.54 cm in 1 inch
• Inexact numbers: those whose values have some uncertainty associated with them.
• Any quantity measured in a lab will be considered inexact.

Uncertainty in Measurements and
Significant Figures
• Uncertainty exists in measurements due to equipment errors (limitations or mis-
calibration within the equipment used to determine the measurement) and
human error (differences in how individuals make a measurement or errors in how
a person interprets a measurement).
• Therefore there are no exact numbers in measurements.The last digit of a
measurement is always uncertain.
• You still want to keep track of the uncertain digit as that gives the precision of the
measurement.
• All digits within a measurement is considered significant, including the uncertain digit.
These are called significant figures.
• The greater the number of digits within the measurement, the greater the precision of
the measurement.

Example Measurements
• Record all certain digits and one uncertain digit when recording measurements
that are not given as a readout (like a digital scale, at which point you would record
all values shown.
• All measurements must include a unit!
cm cm
mL
1.8 cm 3.42 cm
21.5 mL

Rules for Determining Significant
Figures in Reported Measurements
• When a value is given instead of recorded, you need to evaluate the
number in order to determine which digits are significant. In a
recorded values, all non-zero digits are significant. Some zeroes are
significant, some are insignificant.
• When trying to determine which zeroes are significant, use the
following rules:
• Leading zeroes are insignificant
• Leading zeroes are zeroes at the beginning of a number and are considered
place-holder zeroes.
• Trailing zeroes are only significant in numbers that contain a decimal
place.
• Trailing zeroes are zeroes at the end of a number.
• Captive zeroes are always significant.
• Captive zeroes are zeroes between two non-zero digits.

Example Problems
14 significant figures
10.056300034100
10250804900
0.0025930050

Significant Figures in Addition and
Subtraction Calculations
• In mathematical operations involving significant figures (such as calculating the
density of a substance using the mass and volume), the answer is reported in such
a way that it reflects the reliability of the least precise operation. Meaning, an
answer is no more precise that the least precise number used to get the answer.
• Note: exact numbers are considered to have an infinite number of significant
figures.
• When doing addition or subtraction with measured values, the answer should
have the same precision (same number of decimal places) as the least precise
measurement (the one with the fewest decimal places) used in the calculation.

Significant Figures in Multiplication and
Division Calculations
• When doing multiplication or division with measured values, the answer should
have the same number of significant figures as the measured value with the least
number of significant figures.

Multiplication and Division with Addition
and Subtraction
• Please Excuse My Dear Aunt Sally.
→ 𝑥𝑛
→ × ÷ → ±
3.4892
× 5.67 − 2.3 =

Scientific Notation
• A simple way in which scientists handle very small or very big numbers.
• In order to convert from decimal notation (the normal way of viewing a value) to
scientific notation, first you move the decimal so that the number (coefficient) is 1
or greater, but less than 10. Next, you count the number of place values
(exponent) that the decimal moved.The exponent will be positive (moved decimal
left) if the initial number is greater than 1, and negative if it is less than one
(moved decimal right). Plug the number into the formula: coefficient x 10exponent.
• Make sure to keep all significant digits while removing any insignificant zeroes.
6.022 𝑥 1023
Exponent
Coefficient Base

Dimensional Analysis and Conversion
Factors
• Also known as unit conversion.
• Problem solving method that uses the fact that any number or expression can be
multiplied by 1 without changing its value.
• Uses conversion factors to go from one unit to another.
• Conversion factor: a multiplier used to convert the units of one quantity into
another unit without changing the value (ie 1 dozen = 12).
• A conversion factor can be made from any two terms that describe the same quantity or
equivalent amounts.
•
1 𝑑𝑜𝑧𝑒𝑛
12
𝑜𝑟
12
1 𝑑𝑜𝑧𝑒𝑛

Steps of DimensionalAnalysis Problems
• Determine the starting and desired units.
• Come up with conversion factors that contain your starting and ending units (as
well as any that might be needed in between).
• Use your initial value and a starting conversion factor to convert your units. Keep
using conversion factors until you get to your desired units.
• Multiply/divide the whole thing out.
• 𝑔𝑖𝑣𝑒𝑛 𝑢𝑛𝑖𝑡 𝑥
𝑑𝑒𝑠𝑖𝑟𝑒𝑑 𝑢𝑛𝑖𝑡
𝑔𝑖𝑣𝑒𝑛 𝑢𝑛𝑖𝑡
= 𝑑𝑒𝑠𝑖𝑟𝑒𝑑 𝑢𝑛𝑖𝑡
conversion
factor

Metric – English Conversions and Metric –
Metric Conversions
1 hr = 60 min 1 min = 60 sec
24 hrs = 1 day 1 kg = 2.2 lbs
1 mi = 5,280 ft 1 kg = 1000 g
365 days = 1 yr 52 weeks = 1 yr
0.621 mi = 1.00 km 1 yd = 36 inches
Prefix Decimal Equivalent
tera- 1,000,000,000,000 .
giga- 1,000,000,000 .
mega- 1,000,000 .
kilo- 1,000 .
hecto- 100 .
deka 10 .
Base unit 1
deci- 0.1
centi- 0.01
milli- 0.001
micro- 0.000001
nano- 0.000000001
pico- 0.000000000001
1 ton = 2000 lbs 7 days = 1 week
1 gal = 3.79 L 264.2 gal = 1 m3
1 lb = 16 oz 20 drops = 1 mL
2.54 cm = 1 in 1 L = 1000 mL
1 cc is 1 cm3
1 mL = 1 cm3

9,474 mm to cm
1 hr = 60 min 1 min = 60 sec 1 ton = 2000 lbs 7 days = 1 week
24 hrs = 1 day 1 kg = 2.2 lbs 1 gal = 3.79 L 264.2 gal = 1 cubic meter
1 mi = 5,280 ft 1 kg = 1000 g 1 lb = 16 oz 20 drops = 1 mL
365 days = 1 yr 52 weeks = 1 yr 2.54 cm = 1 in 1 L = 1000 mL
0.621 mi = 1.00 km 1 yd = 36 inches 1 cc is 1 cm3
1 mL = 1 cm3
9474 𝑚𝑚
1 𝑚
1000 𝑚𝑚
100 𝑐𝑚
1 𝑚
= 947.4 𝑐𝑚
9474 𝑚𝑚
1 𝑐𝑚
10 𝑚𝑚
= 947.4 𝑐𝑚

32
𝑓𝑡
𝑠𝑒𝑐
to
𝑚𝑒𝑡𝑒𝑟𝑠
𝑚𝑖𝑛
1 mL = 1 cm3
32
𝑓𝑡
𝑠𝑒𝑐
12 𝑖𝑛𝑐ℎ𝑒𝑠
1 𝑓𝑡
2.54 𝑐𝑚
1 𝑖𝑛𝑐ℎ
1 𝑚
100 𝑐𝑚
60 𝑠𝑒𝑐
1 𝑚𝑖𝑛
= 585.216
𝑚
𝑚𝑖𝑛
= 590
𝑚
𝑚𝑖𝑛

If a person weighs 125 lbs, 8 oz., how
many mg does s/he weigh?
1 mL = 1 cm3
8 𝑜𝑢𝑛𝑐𝑒𝑠
1 𝑙𝑏𝑠
16 𝑜𝑢𝑛𝑐𝑒𝑠
= 0.5 𝑙𝑏𝑠 + 125 𝑙𝑏𝑠 = 125.5 𝑙𝑏𝑠
1 𝑘𝑔
2.2 𝑙𝑏𝑠
1000 𝑔
1 𝑘𝑔
1000 𝑚𝑔
1 𝑔
= 5.7 x 107
mg

Using Density as a Conversion Factor
• Can be used to convert mass to volume or volume to mass.
• Density is generally given in g/cm3 or g/mL.
• SI unit for density is kg/m3.
• Density of pure water is 1.0 g/mL or g/cm3.

What is the Mass in Grams of 5.80 Gallons
ofWater?
5.80 𝑔𝑎𝑙𝑙𝑜𝑛𝑠
3.79 𝐿
1 𝑔𝑎𝑙𝑙𝑜𝑛
1000 𝑚𝐿
1 𝐿
1 𝑔
1 𝑚𝐿
= 21982 𝑔 = 2.20 𝑥 104
𝑔
Density
of water
1 mL = 1 cm3

Separation of Mixtures using Differences
in Properties
• Using differences in properties (ie differences in size, mass, melting point, boiling
point, etc) a mixture can be broken down into the component pure substances.
• Some common separation techniques include filtration (separation by differences
in physical state), distillation (separation by differences in melting points), and
chromatography (separation by differences the ability of a substance to adhere to
a solid substrate)

Filtration
• Used to separate a solid from a fluid (liquid or gas)
• Achieved by using a medium (filter) which allows the fluid to pass
through, while trapping the solid.
• The fluid that passes through is called the filtrate.
• There are two common filtration techniques, gravity filtration and
vacuum filtration.
• In gravity filtration, you want to isolate the filtrate (removing the solid
impurities).
• Uses a regular funnel and assisted by gravity.
• In vacuum filtration, you want to isolate the solid from the fluid
(drying the solid in the process).
• Uses a Buchner funnel and assisted by a vacuum.

Distillation
• Separation technique depends on
the ability of a substance to form a
gas.The components must have
sufficiently different boiling points
and heats of vaporization.
• The mixture is heated up, one
substance enters the vapor phase,
then is drawn to another area to be
cooled and recondensed.

The Scientific Method
• A process of experimentation that is used to explore
observations and answer questions.
• Given in a series of steps
• A means scientists use to gather and interpret data
(measurements).

Hypothesis,Theory, and Scientific Law
• Hypothesis = a proposed (testable) explanation for an observed phenomenon
made on the basis of limited evidence as a starting point for further investigation.
• Theory = a hypothesis that has been “proven” with enough data to support it. It is
a model with predictive powers
• Scientific law = when nature behaves in a certain way over and over again, under
multiple different conditions (Law of definite proportions).

The Steps of the Scientific Method

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