Download Types of Measurement

GHSGT
Chemistry Review
Types of Measurement
Qualitative measurement: uses words to
describe something. (i.e. A yellow-green
gas was released.)
Quantitative measurement: uses numbers
to describe something. (i.e. The oxide has
a mass of 1.567 grams.)
Temperature Scales
Celsius, Kelvin, Fahrenheit
Converting Celsius to Kelvin:
C + 273 = ____ K
a. Example: 25 C = ______
25 + 273 = 298 K
Converting Kelvin to Celsius:
K – 273 = ______ C
a. Example: 300 K = _________ C
300 – 273 = 27 C
Classification of Matter
The law of conservation of matter states that
matter can neither be created nor destroyed.
anything that has mass
Matter
and occupies space
Are the properties and
composition constant?
no
mixture
physical change
seawater
mud
fruit salad
yes
yes
pure substance
Is chemical separation
into simpler substances
possible?
no
chemical compound
element
water, sodium chloride
oxygen, gold,
sulfur
Classification of Matter
Compound: A substance with a constant
composition that can be broken down into
elements by chemical processes.
Element: A substance that cannot be broken
down into simpler substances by chemical
means.
Classification of Matter
Mixture
1. Components of two or more components that
can be separated by physical means, such
as chromatography, filtration, or distillation
2. No chemical bonds between the separate
parts
3. Components can vary in composition.
Classification of Matter
Two kinds of mixtures
1. Homogeneous mixtures (solutions)
A. can contain solids, liquids, or gases
B. are evenly mixed, appear as one
component even though there are two or
more parts.
1. Examples: plain jello, air, and salt water
Classification of Matter
2. Heterogeneous mixtures
A. are unevenly mixed
B. can see separate components
C. Examples: dirt, concrete, salad dressing
Classification of Matter
Classify these as an element, a compound, a
heterogeneous mixture, or a homogeneous
mixture.
1. Concrete (heterogeneous mixture)
2. Air (homogeneous mixture)
3. Salt (compound)
4. Gold (element)
5. Helium (element)
6. Tea (homogeneous mixture)
7. Salt water (homogeneous mixture)
Metric Prefixes
Metric Prefixes and Powers of 10
1. mega = 106 (megagrams) (Mg)
2. kilo = 103 (kilograms) (kg)
3. base = 100 (grams) (g)
4. centi = 10-2 (centimeters) (cm)
5. milli = 10-3 (milligrams) (mg)
6. micro = 10-6 (micrograms) (μg)
6000 mg = 6 g = 0.006 kg
Precision vs. Accuracy
Precision refers to how close a series of
measurements are to one another.
1. Replication of result / value
Accuracy refers to how close a measured value
is to an accepted value.
1. True or correct result / value
Seven Base SI Units
1. Length
2. Mass
3. Time
4. Temperature
5. Current
6. Amount
7. Luminousity
meter
(m)
kilogram (kg)
second
(s)
Kelvin
(K)
Ampere (amp)
mole (mol)
candela (cd)
Derived Units
Derived Units are combinations of two or more
units.
1. Examples: speed = mi/hr, ft/s, m/s
density = mass/volume
area = (length) (width)
Density
Density is the mass of a substance per unit
volume of the substance.
1. density = mass/volume (g/ml or g/cm3)
2. Water (the standard for all density values) =
1.0 g/ml
3. 1ml = 1cm3
4. What is the volume of a 50 g metal block
with a density of 5 g/cm3 ?
5. d = m/V, m = dV, V = m/d
The Chemists’ Shorthand:
Atomic Symbols
mass number =
atomic number =
39
K
19
= element
symbol
Atomic Structure
atomic number = the number of protons
1. The atomic number also represents number
of electrons in an neutral atom
mass number = the total number of protons
plus neutrons in the nucleus
Atomic Structure
In uncharged (neutral) atoms, the atomic
number equals the number of protons equals
the number of electrons.
If an atom is charged (different number of
protons and electrons), then it is called an ion.
1. A (+) charged ion is called a cation.
2. A (–) charged ion is called an anion.
Atomic Structure
Isotopes are atoms of the same element (with
the same number of protons) but different
number of neutrons.
1. An element is defined by the number of
protons in the atom’s nucleus.
A. Example: carbon – 12 (12 neutrons);
carbon – 14 (14 neutrons)
Periodic Table of the Elements
Atomic Structure
Atomic Structure
Atomic Structure
Particle
Mass
Charge
Electron
9.1094 x 10-31 kg
-1
Proton
1.6726 x 10-27 kg
+1
Neutron
1.6749 x 10-27 kg
0
Periodic Table of the Elements
Groups (vertical - up and down the table)
1A = alkali metals
2A = alkaline earth metals
7A = halogens
8A = noble gases
Periods (horizontal - across the table)
numbered 1-7
Periodic Table of the Elements
Periodic Table of the Elements
Metals vs Nonmetals - see staircase on the
Periodic table
1. Metals - to the left of the staircase; mostly
solids; conduct electricity; lose electrons to
form positive ions
2. Nonmetals - to the right of the staircase;
most are gases; nonconductors of electricity;
gain electrons to form negative ions
3. Metalloids - border the staircase; have
properties of metals and nonmetals.
Periodic Table of the Elements
Periodic Table of the Elements
There are 7 periods on the periodic table
numbered 1 - 7.
1. They represent the major energy levels (n).
2. They are horizontal rows that extend from
left to right.
Ex: Period 2 includes Li - Ne.
Periodic Table of the Elements
Periodic Table of the Elements
Groups
1. IA - Alkali Metals (1 valence electron) very
reactive
2. IIA - Alkaline Earth Metals (2 valence
electrons)
3. VIIA - Halogens (7 valence electrons) very
reactive
4. VIIIA – Noble Gases (8 valence electrons
except for Helium) non-reactive (very stable)
Periodic Table of the Elements
Periodic Table of the Elements
Representative Elements - The Group “A”
Elements which include all the Groups IA to
VIIIA
Transition Elements - The Group “B” Elements
Periodic Table of the Elements
Periodic Table of the Elements
Lanthanide Series - the 4f row that includes
# 57 (Lanthanum) through # 71 Lu
Actinide Series - the 5f row that includes # 89
Ac (Actinum) through # 102 No
Periodic Table of the Elements
Periodic Table of the Elements
Atomic Radius - the distance from the center of
the nucleus to the outermost valence shell
Periodic Trend - The atomic radius increases
as one moves down the group. The atomic
radius decreases as one moves across a
period.
Periodic Table of the Elements
Inorganic Compound Classification
Two main kinds of compounds
1. Ionic: made up of ions of opposite charge
A. strong electrostatic force of attraction;
ionic bond
B. electrons are transferred
2. Covalent: made up of two or more nonmetals
A. electrons are shared
Chemical Bonds
Atoms bond with each other to become more
chemically stable than they were before they
bonded. To do this their outer electron
(valence) shell must be complete.
1. The Octet Rule states that atoms will either
gain, lose , or share valence electrons to
attain “8” electrons in their outer (valence)
shell to become stable.
2. The Noble Gases are the only one group of
elements that are already stable.
Naming Binary Ionic Compounds
Ions of opposite charge are bonded together.
1. The metal cation (+) is written first and is
named by the metal’s name
2. The nonmetal anion (-) is written second and
is named by the nonmetal’s name with a
revised ending of - ide.
3. Net charge of ions in the compound = 0.
4. Subscripts are used to indicate the number
of ions needed to attain the necessary net
charge of 0.
Naming Binary Ionic Compounds
Examples of Binary Ionic Compounds
What would the formulas be?
1. Sodium chloride
2. Lithium nitride
3. Calcium Fluoride
What would the names be?
1. Al2S3
2. BaO
3. MgBr2
Periodic Table of the Elements
Covalent (Molecular) Compounds
Made up of nonmetals that share
electrons between atoms.
1. This type of bond is called a
covalent bond.
Naming Binary Covalent
Compounds
1. The first nonmetal’s name is that of the
element.
2. The second nonmetal’s name has an -ide
ending, just like with ionic compounds.
3. Use prefixes to describe the subscripts
(1 - mono; 2 - di; 3 - tri; 4 - tetra; 5 - penta;
6 - hexa; 7- hepta; 8 - octa; 9 - nona;
10 - deca)
Examples of Covalent Compounds
P2O5 - diphosphorus pentoxide
How would you write the following formulas?
1. Carbon monoxide
2. Tetranitrogen decoxide
3. Dinitrogen hexoxide
Balancing Chemical Equations
The Law of Conservation of Matter
1. The mass of the products = the mass of the
reactants (matter is not created or
destroyed).
2. Chemical equations are balanced to ensure
that a chemical reaction follows the law of
conservation of matter.
Five Balancing Equation
Guidelines
1. Count the number of atoms of each element
on both the reactant and the product side.
2. Use coefficients (the numbers in front of the
chemical symbol or formula)
3. Never add or change the subscripts.
4. There are Seven Diatomic Elements
(N2, O2, F2, Cl2, Br2, I2, H2)
5. Balance the hydrogen atoms and the
oxygen atoms last.
Balancing Chemical Equations
6. A chemical equation uses symbolic language
to describe a chemical reaction.
7. Equation means equal numbers of atoms of
each element on both sides.
8. Quantities of reactants and products are
expressed in moles by using coefficients.
Balancing Chemical Equations
Symbols for balancing equations
(s) – solid; (g) – gas; (l) – liquid
(aq) – aqueous (dissolved in water)
Go to www.usaprep.com to practice balancing
equations.
Balancing Chemical Equations
Try these equations:
1. H2 + O2
H 2O
2. Ca +
Br2
3. Ba(s) +
4. Mg +
CaBr2
O2(g)
AuCl3
BaO(s)
MgCl2 +
Au
Balancing Chemical Equations
The Answers:
1. 2H2 + O2
2H2O
2. Ca + Br2
CaBr2
3. 2Ba(s) + O2(g)
2BaO(s)
4. 3Mg + 2AuCl3
3MgCl2 + 2Au
Acids and Bases
Acid Definitions:
1. Sour taste
2. Neutralize the actions of bases
3. Blue litmus paper turns red
4. Liberates hydrogen gas when reacted with
certain metals
5. Examples: Foods and Drinks
Acids and Bases
Base Definitions:
1. Bitter taste
2. Neutralizes the action of acids
3. Slippery to the touch
4. Red litmus paper turns blue
5. Examples: Cleaning Solutions
Acids and Bases
an acid + a base
a salt + water
Strong acids:
hydrochloric acid, sulfuric acid, nitric acid
Strong base:
sodium hydroxide
Weak acid:
acetic acid
Weak base:
ammonia
pH Scale
The pH scale has
values from 0 - 14.
0 - 6 is acidic
7 is neutral
8 -14 is basic
The Mole
The mole is the SI unit of measure for the
amount of a substance.
1. The mole is a way to measure the mass of
elements and compounds.
2. The molar mass of any element is
numerically equal to its atomic mass and has
the units of g/mol.
3. Example: The mass of one mole of
potassium (K) is 40 grams.
Physical and Chemical Changes
1. Physical Change - Only appearance changes.
Identity still the same. Kinds of changes are:
Size/Shape and Phase (State).
A. Size and Shape (splitting, breaking,
tearing, hammering, etc.).
B. Phase (State) (melting, vaporizing,
freezing, condensing, etc.).
Physical and Chemical Changes
2. Chemical Change - The appearance and the
identity changes. A new product is formed.
The way you know a new product is formed
is by the following:
A. gas formed
B. color change
C. mass change
D. heat change
E. solid formed
F. light released
Examples are: rusting, growing, burning, cooking,
combusting, fermenting, frying, and exploding.
Nuclear Reactions
A nuclear reaction results when an unstable nucleus
breaks down and emits radioactive particles.
1. There are 3 types of particles released during
radioactive decay.
A. alpha particle (α) released during alpha decay
B. beta particle (β) released during beta decay
C. gamma ray (γ) released during any decay
Three Types of Radioactive Decay
1. alpha particle (helium nucleus (+ 2 charge)
A. largest, slowest, least penetrating particle
2. beta particle (fast electron (– 1 charge)
A. more penetrating than an alpha particle
3. Gamma radiation (no particle, no mass or
charge – pure energy)
A. most penetrating and most damaging
(shielded by lead)
Half Life
The half life of a radioactive isotope is the time
it takes for one half of the isotope to decay.
1. Example: The half life of mercury - 195 is 31
hours. If you start with 20 g, how much will
be left after
(A) 31 hours? (10 g)
(B) 62 hours? (5 g)
Particles in an Electric Field
Remember the charges for each particle. Also
remember that like charges repel and opposite
charges attract.
Nuclear Fission and Nuclear Fusion
Nuclear Fission is the splitting of large nuclei
resulting in a tremendous release of energy.
Nuclear Fusion is the combining of small nuclei
resulting in an even greater release of energy.
The sun uses nuclear fusion to produce energy.
1. 2 hydrogen atoms combine to form 1 helium
atom.
Energy/ Heat/ Phase Changes
1. Temperature – the measure of the average
kinetic energy of particles.
2. Heat – form of energy that may be absorbed
or released. Flows from a warm body to a
cooler one until equilibrium is reached.
3. A calorie – the amount of heat required to
raise the temperature of one g of water one
degree Celsius.
Types of Phase Changes
1. melting - solid to liquid
2. freezing - liquid to solid
3. evaporation - liquid to gas
4. condensation - gas to liquid
5. sublimation - solid to gas
6. deposition - gas to solid
Key Terms for Heat
1. Energy – the capacity to do work
2. Heat – energy transferred from one object to
another (The SI unit of heat is the Joule (4.18
Joules = 1 calorie).
3. Thermochemistry – the study of heat effects
in chemical reactions
4. Combustion – chemical reactions that
release heat
5. Exothermic reaction – one that releases heat
6. Endothermic reaction – one that absorbs heat
Heat
Specific heat – amount of heat required to raise
the temperature of one gram of a substance 1o
Celsius. Specific heat of water is 1.
1. Water – high specific heat due to hydrogen
bonds. Water doesn’t change temperature
very much despite large amounts of heat
energy added or subtracted.
2. Metals have low specific heat capacities.
Energy in Chemical Reactions
The bond breaking that occurs in reactants
during a chemical reaction requires energy.
The bond formation that occurs in products
during a chemical reaction releases energy.
An endothermic or exothermic reaction is
determined by the balance between these two
processes.
3 Phases (States) of Matter
The balance between the attractive forces
between the particles and the kinetic energy of
the particles determines the phase of matter.
1. High KE of the particles and low attractive
forces between the particles = gas.
2. Low KE of the particles and high attractive
forces between the particles = solid.
3. Intermediate KE of the particles and
intermediate attractive forces between the
particles = liquid.
The Phase Diagram for Water
Solids
The two main types of solids are crystalline
solids and amorphous solids.
1. In a crystalline solid, the atoms, ions, or
molecules are arranged in an orderly,
repeating, 3-dimensional pattern (crystal
lattice).
Solids
2. In an amorphous solid, the internal structure
lacks order. Atoms, ions, or molecules are
randomly arranged.
A. These substances usually cool rapidly –
Not enough time for the particles to
arrange themselves in a pattern.
B. Examples: Rubber, glass, plastics,
polymers
Allotropes
Allotropes – substances with same elemental
composition, different geometric arrangements.
1. Example, carbon has 4 allotropes:
A. diamond - formed under tremendous
pressure
B. graphite - more loosely packed
C. soot - randomly bonded (amorphous
form)
D. buckey ball
Kinetic Molecular Theory of Gases
1. Gases – made up of very small particles
(atoms or molecules) and separated by large
distances (low density).
2. Particles are in constant, random, straight-line
motion – thousands of perfectly elastic
collisions per second. Total KE remains
constant.
3. Gases exert a pressure due to the collisions
on each other.
Collision Theory
1. When gas particles collide they exert a
pressure on their container.
2. Temperature is the measure of the average
kinetic energy of the gas particles.
3. There are four main properties of gases that
determine their physical behavior.
Four Main Properties of Gases
1. Pressure = force / area
2. Temperature = the average kinetic energy of
particles
3. Volume = space occupied by matter.
4. Amount of gas is measured in grams or
moles.
What happens to the volume of a gas if the
pressure is increased at constant temperature?
What happens to the volume of a gas if the
temperature is increased at constant pressure?
Lewis Dot Structures
Rules for Lewis Dot Structures
1. Count total number of dots (valence electrons)
in the structure.
2. Spatially arrange the atoms. (More than two
atoms – find the central atom)
3. Obtain 8 dots around each atom (2 around
hydrogen).
4. If single bonds don’t work, try double bonds,
then triple bonds.
Three Factors Affecting the
Rate of Dissolving
How fast will a solute dissolve in a solvent?
1. Stirring or agitation (more solute / solvent
contact at a faster rate)
2. Smaller particles (increases the surface area
of the solute, therefore there is more solute /
solvent contact at a faster rate)
3. Increasing the temperature (increases the
kinetic energy and faster rate of contact
between the solute/solvent particles)
The Solubility of Solids at
Different Temperatures
1. Solubility of sugar at
50oC?
2. Which solute least
affected by increase
in temp?
3. Generally, as temp
increases, solubility?
GHSGT Chemistry Review
The End