Download 9/6/12

9/6/12
What is chemistry?
A chemical is any substance that has a definite composition.
A chemical reaction is the process by which one or more substances change to
produce one or more new substances.
Physical states of matter
-The states of matter are the physical forms of matter which are solid, liquid, gas,
and plasma.
Macroscopic refers to what you can see with the unaided eye
Microscopic refers to what you would see if you could see individual atoms.
9/7/12
Solids have a fixed volume and shape that result from the way their particles are
arranged.
Liquids have a fixed volume but not a fixed shape
Gases have neither a fixed volume nor shape
Physical Changes are changes in which the identity of a substance doesn’t change
- Changes of state are physical changes
Chemical Changes occur when the identities of substances change and new substances
form.
9/10/12
Chemical Changes
Mercury(II) oxide  mercury + oxygen
Reactants are the substances on the left-hand side of the arrow
(they are used up in the reaction
Products are the substances on the right-hand side of the arrow
(they are produced in the reaction)
Evidence of a chemical change
- the evolution of a gas
- the formation of a precipitate (solid substance forms and typically sinks to the
bottom)
- the release or absorption of energy
- a change in temperature or the giving off of light
- a color change in the reaction system
Matter
Matter- Peanut butter, water, fish, garbage, human brain, carbon dioxide, air,
yourself, tree
Not Matter- light, time, motion, an idea, energy
Distinguish between different characteristics of matter, including mass, volume, and
weight.
Identity and use SI units in measurements and calculations
Set up conversion factors, and use them in calculations
Identify and describe physical properties, including density
Identify chemical properties
Matter had Mass and Volume
- Matter is anything that has mass and volume
- Volume is the space na object occupies
- Mass is the quantity of matter in an object
o Devices used for measuring mass in a laboratory are called balances or
scales
- Weight is the force produced by gravity acting on a mass
9/11/12
Units of Measurements
- when working with numbers, be careful to distinguish a quantity and its units
o Quantity describes something that has magnitude, size, or amount
o Unit is a quantity adopted as a standard of measurement
Meter
kilogram
second
ampere
Unit
symbol
m
kg
s
A
kelvin
K
candela
cd
mole
mol
Unit name
Name
Symbol
Factor
Prefix
exa
peta
tera
giga
mega
kilo
hecto
deca
Symbol
E
P
T
G
M
k
h
da
deci
centi
milli
micro
nano
pico
d
c
m
μ
n
p
100
1000m
10006
10005
10004
10003
10002
10001
10002/3
10001/3
10000
1000-1/3
1000-2/3
1000-1
1000-2
1000-3
1000-4
decada
101
10n
1018
1015
1012
109
106
103
102
101
100
10-1
10-2
10-3
10-6
10-9
10-12
Quantity name
length
mass
time
electric current
thermodynamic
temperature
Dimension
symbol
L
M
T
I
Quantity symbol
l (a lowercase L), x, r
m
t
I (an uppercase i)
T
Θ
luminous intensity
Iv (an uppercase i with lowercase
non-italicized v subscript)
J
amount of substance
n
N
hectoh
102
kilo- megak
M
103
106
giga- teraG
T
109
1012
peta- exa- zettaP
E
Z
1015
1018 1021
Decimal
Short scale
1000000000000000000
quintillion
1000000000000000 quadrillion
1000000000000
trillion
1000000000
billion
1000000
1000
100
10
1
0. .1
0. .01
0. .001
0. .000001
0. .000000001
0 .000000000001
billionth
trillionth
Long scale
trillion
billiard
billion
milliard
million
thousand
hundred
ten
one
tenth
hundredth
thousandth
millionth
milliardth
billionth
A conversion factor is a simple ration that relates two units that express a measurement of
the same quantity
you can construct conversion factors between kg and grams as follows:
1kg/1000g or 1000g/1kg
Given: 4.5kg
?g
4.5kg 1000g = 4500g
1kg
-
Given: 208lb
? kg
208lb .45kg = 93.6kg
1lb
9/13/12
Given: 0.851L
?mL
.851L 1000mL = 851.0mL
1L
Derived Units
- Many quantities you can measure need units other than the seven basic SI
units
- These units are derived by multiplying or dividing the base unit
o Speed is distance divided by time. A derived unit of speed is
meters/second (m/s)
o A rectangle’s area is found by multiplying its length (in meters) by its
width (also in meters)
 Its unit is square meters (m2)
- Volume is another commonly used derived unit
o The volume of a book can be found by multiplying its length, width,
and height.
o The unit of volume is the cubic meters (m3)
o This unit is too large and inconvenient in most labs. Chemists usually
use the liter (L)
 1L = 1000mL = 1000cm3
 1mL = 1cm3
Properties of Matter
Physical Properties
- A physical property of a substance is a characteristic that does not involve a
chemical change.
- Physical properties of a substance can be determined without changing the
nature of a substance.
- Physical properties include malleability, color, texture, state, density, melting
point, and boiling point.
9/14/12
Properties of Matter
Density is the Ratio of Mass to Volume
-
The density of an object is the mass if the object divided by volume of the
object
Densities are expressed in derived unties such as g/cm3 or g/mL
Density is calculated as follows:
o Density = mass/volume OR D = m/v
The density of a substance is the same no matter what the size of the sample
is.
Because the density of a substance is the same for all samples, you can use
this property to help identify substances.
9/18/12
Do Now
Kilo = 1000
.01 = centi
45g = .045 kg
100mL = .1L
4.8cg = 48mg
Properties of Matter
Chemical Properties
- A chemical property is a property of matter that describes a substance’s ability
to participate in chemical reactions – i.e.- Reactivity
- A chemical property of many substances is their reactivity with oxygen.
o Rusting, corrosion
- Some substances break down into new substances when heated
Classifying Matter
- An atom is the smallest unit of an element that maintains the properties of that
element.
- Matter exists in many different forms but there are only 118+ types of atoms.
- Atoms are joined together to make up all the different kinds of matter.
Pure Substance
- A pure substance is a sample of matter, either a single element or a single
compound, which has definite chemical and physical properties.
- Elements are pure substances that only contain one kind of matter (atom).
They cannot be separated or broken down into simpler substances by ordinary
chemical means. (exception- Higgs Boson particle)
o Each element has its own unique set of physical and chemical
properties
- A molecule is the smallest unit of a substance that keeps all of the physical
and chemical properties of that substance.
- A molecule usually consists of two or more atoms combined in a definite ratio
o Except diatomic elements
- Diatomic elements exist as two atoms of the same element joined together.
o H2 O2 N2 F2 Cl2 Br2 I2
o Acronym- HOFBrINCl
- Some elements, such as oxygen, phosporus, sulfur, and carbon, have many
molecular forms
- An allotrope is one of a number of molecular forms of an element
-
The properties of allotropes vary widely
o Carbon- graphite, diamond, buckyballs
9/20/12
Pure Substances
Some Elements Exist in More than One Form
- Oxygen exists as allotropes
- Oxygen gas (O2) is colorless and odorless
- Ozone (O3) is toxic and pale blue
Compounds are Pure Substances
- Pure substances that are not elements are compounds. Compounds are
composed of more than one kind of atom.
o Example: carbon dioxide
- There may be easier ways of preparing them, but compounds can be made
from their elements.
- Compounds can be broken down into their elements, often with difficulty.
Compounds are Represented by Formulas
- Compounds have characteristic in properties and composition
- Compounds can be represented by an abbreviation or Formula
o A formula has subscripts, which represent the number of different
atoms in the compound.
o Example: H2O has 2 hydrogen atoms and 1 oxygen atoms
- Molecular formulas give information only about what makes up a compound.
o Example: the molecular formula for aspirin is C9H8O4
- A structural formula shows how the atoms are connected
o Two-dimensional models do not show the molecule’s true shape
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How is matter classified?
- A call and stick model shows the distances between atoms and the angles
between them in three dimensions.
- A space filling model attempts to represent the actual size of the atoms nd not
just their relative position
- These models convey different information
Mixtures
- A mixture is a combination of two or more substances that are not chemically
combined
- Air is a mixture of mostly nitrogen and oxygen
- Water is not a mixture
o The H and O atoms are chemically bonded
o The ratio of H to O atoms is always 2 to 1
- The properties of the mixture may vary
- An alloy is a solid mixture
o Example: an alloy of gold and other metal atoms is stronger than pure
gold
 18-karot gold contains 18 grams of gold per 24 grams of alloy
 14-karot gold contains 14 grams of gold per 24 grams of alloy
Homogenous mixture
- A homogenous mixture describes something that has a uniform structure or
composition throughout.
o Example: gasoline, syrup, and air
- Homogenous mixtures have the same properties throughout
Heterogeneous Mixtures
- A heterogeneous mixture describes something that is composed of dissimilar
components
o Example: a mixture of sand and water
- Any two samples of a heterogeneous mixture will have the different
proportions of ingredients
o Heterogeneous mixtures have different properties throughout
- Mixtures can be either homogenous or heterogeneous
9/24/12
Do Now
37 gigabytes = 37,000,000 kilobytes
763 mL = .763L
45 kilograms = 45,000,000 mg
Understanding Concepts
Tea is best classified as a homogenous mixture
He is an element
CHAPTER 2
Objectives
- Explain that physical and chemical changes in matter involve transfers of
energy
- Apply the law of conservation of energy to analyze changes in matter
- Distinguish between heat and temperature
- Convert between Celsius and Kelvin
Energy and Change
- Energy is the capacity to do work, such as moving an object, forming a new
compound, or generating light (Electromagnetic Energy)
- Energy is always involved when there is a change of matter
Changes in Matter Can Be Physical or Chemical
- Ice melting and water boiling are examples of physical changes
- A Physical Change is a change of matter that affects only the physical
properties of matter.
- In contrast, the reaction of hydrogen and oxygen to produce water is an
example of a chemical change
- A Chemical Change is a change that occurs when one or more substances
change into entirely new substances with different properties
- A chemical change occurs whenever a new substance is made
- All physical and chemical changes involve a change in energy
- Sometimes energy much be supplied for the change in matter to occur
- Evaporation is the change of a substance form a liquid to a gas due to the
absorption of Energy
9/25/12
Energy and Change
Endothermic and Exothermic Processes
- Any change in matter in which energy is absorbed from the surrounding in an
Endothermic process
o The melting of ice and boiling of water are examples of physical
changes that are Endothermic
o As the chemicals react, energy is absorbed. Energy is a Reactant
- Any change in matter in which energy is released is an Exothermic process
o The freezing of water and condensation of water vapor are two
examples of physical changes that are exothermic processes.
o Energy is released. Energy is a Product.
- Energy can be absorbed (Endothermic) by the surroundings or released
(exothermic) to the surroundings, but it cannot be created or destroyed.
- The Law of Conservation of Energy states that during any physical or
chemical change, the total quantity of energy remains constant.
- In other words, energy cannot be created nor destroyed.
Energy is often transferred
- A system consists of all the components that are being studied at any given
time.
o The chemicals are the system, not the beaker they are in.
- The surroundings include everything around outside the system.
- Energy exists in different forms, including:
o Heat, Light, Potential-Chemical, kinetic-Motion, Sound, Electrical,
Pressure waves
- The transfer of energy between a system and its surroundings can involve any
one of these forms of energy
Heat
- Heat is the energy transferred between objects that are different temperatures.
- Heat energy is always transferred from a warmer object to a cooler object
- For example, when ice cubes are placed in water, heat energy is transferred
from the water to the ice.
- Energy is also transferred as heat during chemical changes.
9/27/12
Energy Can Be Absorbed as Heat
- In an endothermic reaction, energy is absorbed by the chemicals that are
reacting.
Heat is Different from Temperature
- Scientists define Temperature as a measurement of the average kinetic energy
of the random motion of the particles in a substance
- The transfer of energy as Heat can be measured by calculated changes in
Temperature
Temperature is Expressed Using Different Scales
- Thermometers are usually marked with the Fahrenheit or Celsius temperature
scales
- A third temperature scale uses the unit Kelvin, K
The zero point of the Celsius scale is designed as the freezing point of water
The zero point on the Kelvin scale is designated as absolute zero, the
temperature at which the minimum average kinetic energies of all particles
occur.
- There are no negative temperatures in Kelvin
Transfer of Heat May Not Affect Temperature
- The transfer of energy as heat does not always result in a change of
temperature
o For example, consider that happens when energy is transferred to a
mixture of ice and water
o As energy is transferred as heat to the ice-water mixture, the ice cubes
start to melt.
o The temperature of the mixture remains at 0C until all the ice has
melted.
-
QuickTime™ and a
decompressor
are needed to see this picture.
Transfer of Heat Affects Substances Differently
- The specific heat of a substance is the quantity of energy as heat that mist be
transferred to raise the temperature of 1g of a substance 1K
- The SI (metric) Unit for energy is Joule (L)
- Specific Heat is expressed in joules per gram Kelvin (J/gK)
- Metals tend to have low specific heats
- Water has a high specific heat
9/28/12
Scientific Method
- The Scientific Method is a series of steps followed to solve problems,
including:
o Collecting data
o Formulating a hypothesis
o Testing the hypothesis
o Stating conclusions
- An experiment is the process by which scientific ideas are tested.
- A Hypothesis is a reasonable and testable explanation for observations.
- A Variable is a factor that could affect the results of an experiment.
- When a variable is kept constant form one experiment to the next, the variable
is called the control.
- The procedure is called a Controlled Experiment
- Any hypothesis that withstands repeated testing may become part of a theory
- In science, a Theory is a well-tested explanation of observations
Theories and Laws Have Different Purposes
- Some facts in science always hold true. These facts are called laws.
- A Law is a statement or mathematical expression that reliably describes a
behavior of the natural world.
- A theory is an attempt to explain the cause of certain events in the natural
world.
- For example, the Law of Conservation of Mass states that the products of a
chemical reaction have the same mass as the reactants have.
- Keep in mind that a hypothesis predicts an event, a theory explains it, and a
law describes it.
Models Can Illustrate the Microscopic World of Chemistry
- A Model represents an object, a system, a process, or an idea
10/3/12
Measurements and Calculations in Chemistry
Accuracy and Precision
- The Accuracy of a measurement is how close the measurement is to the true
or actual value
- Precision is the exactness of a measurement
- It refers to how closely several measurements of the same quantity made in
the same way can agree.
Significant Figures
- Scientists report values using significant figures
- The Significant Figures of a measurement or a calculation consists of all the
digits known with certainty as well as one estimated, or uncertain, digit.
- The last digit or significant figure reported after a measurement is uncertain
and estimated
Rules for Determining Significant Figures
- Nonzero digits are always significant
- Zeros between nonzero digits are significant
- Zeros in front of nonzero digits are not significant
- Zeros both at the end of a number and to the right of a decimal point are
significant
- Zeros both at the end of a number but to the left of a decimal point may not be
significant. If a zero has not been measured or established, it is not significant.
A decimal point placed after zeros indicated that the zeros are significant
10/8/21
Do Now
Define: Density- mass divided my volume
Formula- D=m/v
Units- g/mL or g/cm3
Significant Figures
- To determine the number of significant figures in the answer, you must first
find the number of significant figures in the values used to calculate the
answer.
When multiplying and dividing, the answer cannot have more sigfigs than the
measurement with the smallest number of sigfigs.
- With addition and subtraction, the result can be no more certain than the least
certain number in the calculation.
- Conversion factors and countable values have unlimited numbers of sigfigs
OR it does not limit my calculation.
10/9/12
Significant Figures
Example: A student heats 23.62g of a solid and observes that the temperature
increases from 21.6C to 36.79C. Calculate the temperature increase per
gram of solid.
36.79C - 21.6C = 15.19C = 15.2C
15.2C/23.62g = 0.644C/g
Specific Heat Depends on Various Factors
- Specific heat depends on the nature of the material that is changing
temperature, the mass of the material, and the size of the temperature change.
- Recall that the specific heat is the quantity of energy that must be transferred
as heat to raise the temperature of 1g of a substance by 1K.
- The specific heat (cp) of a substance at a given pressure (p) is calculated by
the following formula
o Specific heat = cp = q / mxT (q-heat in joules, m-mass in grams,
T- change in temperature in K)
Example: A 4.0g sample of glass was heated from 274K to 314K and was found
to absorb 32J of energy as heat. Calculate the specific heat of this glass.
cp = q
= 32J
= .20J/gK
mxT (4.0g)(40K)
10/10/12
Scientific Notation
- Very large and very small numbers are often written in scientific notation.
- A number in scientific notation has 2 parts
- The first part is a number that is between 1 and 10
- The second part consists of a power of 10
- To write the first part of the number, move the decimal to the right of the left
so that only one nonzero digit is to the left of the decimal.
Scientific Notation with SigFigs
- Use scientific notation to eliminate all place-holding zeros
- Move the decimal in an answer so that only one digit is to the left, and change
the exponent accordingly. The final value must contain the correct numner of
sigfigs.
o 1001000000 = 1.001x109
o 0.0000456 = 4.56x10-5
o 0.0000036mm = 3.6x10-6mm
o 1450000mg = 1.45x106mg
-
10/11/12
SigFig Practice
1890.01m - 6 sigfigs
0.01702L – 4 sigfigs
Specific heat- is the quantity of energy that must be transferred as heat to raise the
temperature of 1g of a substance by 1K.
Specific heat = cp = q / mxT (q-heat in joules, m-mass in grams, T- change in
temperature in K)
10/15/12
CHAPTER 3
3.1 Atomic Theory
- The idea of an atoms theory is more than 2000 years old.
- Until recently, scientists had never seen evidence of atoms.
- The law of definite proportions, the law of conservation of mass and the law
of multiple proportions support the current atomic theory
The Law of Definite Proportions
- The Law of Definite Proportions states that a chemical compound always
contains the same elements in exactly the same proportions by weight or
mass.
- The law of definite proportions also states that every molecule of a substance
is made of the same number and types of elements.
The Law of Conservation of Mass
- The law of conservation of mass states that mass cannot be created nor
destroyed in ordinary chemical and physical changes.
- The mass of the reactants is equal to the mass of the products.
The Law of Multiple Proportions
- The law of multiple proportions states that when two elements combine to
form two or more compounds, the mass of one element that combines with a
given mass of the other is in the ratio of small whole numbers.
Dalton’s Atomic Theory
- In 1808, John Dalton developed an atomic theory.
- Dalton believed that a few kinds of atoms made up all matter.
- According to Dalton, elements are composed of only one kind of atom and
compounds are made from two or more kinds of atoms.
- Dalton’s Theory Contains Five Principles
o All matter is composed of extremely small particles called atoms,
which cannot be subdivided, created, or destroyed.
o Atoms of a given element are identical in their physical and chemical
properties.
o Atoms of different elements differ in their physical and chemical
properties.
o Atoms of different elements combine in simple, whole number ratios
to form compounds.
o In chemical reactions, atoms are combined, separated, or rearranged
but never created, destroyed, or changed.
Data gathered since Dalton’s time shows that the first two principles are not
true in all cases.
HW- Read Section 3.1, do 1-9 (pg 78)
10/16/12
3.2 Structure of Atoms
Subatomic Particles
- Experiments by several scientists in the mid-1800’s led to the first change to
Dalton’s atomic theory.
- The smaller parts that make up atoms are called subatomic particles
- The three subatomic particles that are most important for chemistry are the
electron, proton, and neutron.
Electrons were discovered using Cathode Rays
- JJ Thompson studied currents
- Thomson observed a glowing beam that came out of the cathode and struck
the anode and the nearby glass walls of the tube.
o He called these rays cathode rays
o The glass tube Thomson used is known as a cathode-ray tube.
An electron has a Negative Charge
- Thomson’s experiments showed that a cathode ray consists of particles that
have mass and a negative charge
- There particles are called electrons
- An electron is a subatomic particle that has a negative electric charge
- Electrons are negatively charged, but atoms have no charge
o Atoms contain some positive changes that are balanced by negative
charges.
- Electron: e, e-, or –10e ; charge of 1.602x10-19C ; common charge notation –1 ;
mass is 9.109x10-31kg
Rutherford Discovered the Nucleus
- Thomson proposed that the electrons of an atom were embedded in a
positively charged ball of matter. His model of an atom was named the plumpudding model.
- Ernest Rutherford performed the gold foil
experiment, which discovered the plum-pudding
model of an atom.
o A beam of small, positively charged
particles, called alpha particles, was
directed at a thing gold foil.
- Rutherford called the space around the atom, the
nucleus.
10/17/12
- The nucleus is the dense, central portion of the atom.
- The nucleus is made up of protons and neutrons
- The nucleus has all the positive charge, and nearly all of the mass, but only a
very small fraction of the volume of the atom.
-
QuickTime™ and a
decompressor
are needed to see this picture.
Protons
- Protons are the subatomic particles that have a positive charge and that is
found in the nucleus of an atom.
o The number of protons of the nucleus is the atomic number, which
determines the identity of an element.
o Because protons and electrons have equal but opposite charges, a
neutral atom must contain equal numbers of protons and electrons.
- Neutrons are the subatomic particles that have no charge and that are found in
the nucleus of an atom
Charge
Particle
Relative Charge
Mass
Relative mass
Proton
+1.60 x 10-19 C
+1
1.672 x 10-24 g
1 amu
Neutron
neutral
0
1.674 x 10-24 g
1 amu
Atomic Number
- The number of protons that an atom has is known as the atomic number.
o The atomic number is the same for all atoms of an element.
o No two elements can have the same atomic number.
- Atomic numbers are always whole numbers.
- The atomic number also reveals the number of electrons in an element.
o For atoms to be neutral, the number of negatively charged electrons
must equal the number of positively charged protons.
10/18/12
Do Now
1) Who discovered the electron? JJ Thomson
How? Cathode Ray Tube
When? Early 1900’s
2) What did Rutherford determine? Atom is made up of mostly space and there
is dense nucleus that is positively charged in the 1900’s
3) Who developed the first Atomic Theory? Dalton in the early 1800’s
Mass Number is the Number of Particles in the Nucleus
- The mass number is the sum of the number of protons and neutrons of an
atom.
o Mass# = p+ + n0
o Mass of e- = 0 (charge –1)
o Mass of p+ = 1 (charge +1)
o Mass of n0 = 1 (charge 0)
- You can calculate the number of neutrons in an atom by subtracting the
atomic number (the number of protons) from the mass number (the number of
protons and neutrons)
o Atomic# - p+
o Mass# - p+ + n0
o # of neutrons = mass# - atomic#
- Unlike the atomic number, the mass number can vary depending on the
number of neutrons.
Example: a particular atom of neon has a mass number of 20. How many
neutrons does is have?
Mass# = 20
Atomic# = 10 (has 10 protons)
Mass# - Atomic# = 20 – 10 = 10 Neutrons
- Sample Problem A: How many protons, neutrons, and electrons are present in
an atom of copper whose atomic number is 29 and whose mass number is 64.
Copper (Cu)
?p+ = 29
?e- = 29
?n0 = mass – atomic = 64 – 29 = 35
- The atomic number always appears on the lower left side of the symbol.
- Mass numbers are written on the upper left side of the symbol.
1
2
2
2
3
7
5
9
11
1H
1H
3He
4He
6Li
3Li
4Be
5B
5B
+
+
+
+
+
+
+
+
1p 2p 2p 2p
3p 7p 5p 9p 11p+
0n0 1n0 0n0 0n0 0n0 4n0 1n0 4n0 6n0
1e- 1e- 3e- 4e- 6e- 3e- 4e- 5e- 5e10/22/12
Atomic Number and Mass Number
Isotopes
- All atom of an element have the same atomic number and the same number of
protons. Atoms do not necessarily have the same number of neutrons.
- Atoms of the same element that have different numbers of neutrons.
- One standard method of identifying isotopes is to write the mass number with
a hyphen after the name of the element.
o helium-3 or helium-4
- The second method of identifying isotopes shoes the composition of a nucleus
as the isotope’s nuclear symbol.
o 32He or 42He
10/23/12
Do Now
-
Element
Magnesium
Magnesium
Symbol
Mg
Mg
Magnesium
chloride
Magnesium
Potassium
ClMg
KMg
Magnesium
Phosphorus
Magnesium
Aluminum
P Mg
AlMg
Magnesium
U
Mg
Au
Uranium
Magnesium
Gold
Molybdenum Mo
Cesium
Cs
p+
n0
e-
Mass#
12
17
19
15
13
92Mg
79
42
55
12
18
21
16
14
146
118
54
78
12
17
19
15
13
92
79
42
55
24
35
40
31
27
238
197
96
133
Atomic#
12
17
19
15
13
92
79
42
55
10/24/12
3.3 Electron Configuration- Atomic Models
Rutherford’s model proposed electron orbits
- The experiments of Rutherford’s team led to the replacement of the plum
pudding model of the atom with a nuclear model of the atom.
o Rutherford suggested that electrons, like planets obiting the sun,
revolve around the nucleus in circular or elliptical orbits.
o Rutherford’s model could not explain why electrons did not crash into
the nucleus.
- Neils Bohr replaced the Rutherford model of an atom 2 years later.
Bohr’s Model Confines Electrons to Energy Levels
- According to Bohr’s model, electrons can be only certain distances from the
nucleus. Each distance corresponds to a certain quantity of energy that an
electron can have.
o An electron that is as close to the nucleus as it can be is in its lowest
energy level.
o The farther an electron is from the nucleus, the higher the energy level
that the electron occupies.
- The difference in energy between two energy levels is known as a Quantum of
energy
- Rutherford had a planetary model
- Bohr had the quantum energy model
Electrons Act like Both Particles and Waves
- Thomson’s experiments demonstrated that electrons act like particles that
have mass.
- In 1924, Louis de Broglie pointed out that the behavior of electrons according
to Bohr’s model was similar to the behavior of waves.
- De Broglie suggested that electrons could be considered waves confined to the
space around the nucleus.
- The present day model of the atom takes into account both the particle and
wave properties of electrons.
- In this model, electrons are located in Orbitals, regions around a nucleus that
correspond to specific energy levels
o Orbitals are regions where electrons are likely to be found.