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Chapter 2- Atoms
HW-2.10, 2.11, 2.15, 2.17, 2.25,
2.29, 2.45, 2.48, 2.53, 2.54, 2.66,
2.79
Ancient Thinkers
• Pondered what matter was made of
• Democritus and followers believed matter
was composed of tiny particles
– Called particles atoms
– Said their were different types of atoms that
add different properties
• Zeno of Elea believed matter was infinitely
divisible
• Atoms- the basic unit of matter
Classifications of Matter
Matter
Anything that occupies space and has mass
Pure Substances
Fixed composition; cannot
Be further purified
Elements
Cannot be
Subdivided by
Chemical or
Physical means
Compounds
Elements united
In fixed ratios
Mixtures
A combination of two
Or more pure substances
Homogeneous
Matter uniform
Composition
throughout
Heterogeneous
Matter nonuniform
composition
Definitons
• Elements- a substance that consists of identical
atoms
• Compounds- a pure substance made up of two
or more elements in a fixed ratio by mass.
• Formula- gives the ratio of a compounds
constituent elements and identifies each
element by its atomic symbol
• Mixture- a combination of two or more pure
substances
Types of Mixtures
• Homogeneous mixture- composition the same
throughout
• Example: Air
– 78% Nitrogen 21% Oxygen
• Heterogeneous mixture- nonuniform
composition
• Example: Blood
• Note: Mixtures consists of two or more pure
substances with different physical properties.
This properties can be used to separate the
mixture.
Dalton’s Atomic Theory
1. All matter is made up of very tiny, indivisible
particles called atoms
2. All atoms of the same element have the same
chemical properties
3. Compounds are formed by the chemical
combination of two or more different kinds of
atoms
4. A molecule is a tightly bound combination of
two or more atoms that acts as a single unit
Laws
• Law of Conservation of Mass
– Matter can be neither created nor destroyed
(This does not contradict Dalton!!)
• Law of Constant Composition
– Compounds are always made up of elements
in the same proportion by mass
Types of Elements
• Monoatomic Elements- elements that
consist of single atoms not connected to
each other. Ex. He, Ne
• Diatomic Elements- elements that exist
as pairs of atoms, bonded together.
Ex. O2 Cl2 Br2 I2 H2 N2
• Polyatomic Elements- elements that
exist with more than two atoms bonded
together. Ex. O3 P4 S8 Diamond
Subatomic Particles
• Proton- subatomic particle found in the
nucleus of an atom and has a positive
charge
• Electron- subatomic particle found in the
space surrounding the nucleus and has a
negative charge
• Neutron- subatomic particle found in the
nucleus and has no charge
Atom Properties
• Mass Number- the sum of the number of
protons plus the number of neutrons
• Atomic Number- the number of protons in
the nucleus of an element
• Symbol for an atomic nucleus- The
element symbol with the mass number as
a superscript on the left and the atomic
number as a subscript on the left.
Properties of Atoms (cont.)
• Isotopes- atoms with the same number of
protons but different numbers of neutrons
– Properties of isotopes of the same element
are almost identical except for radioactivity
properties.
• Atomic weight- of an element in the
Periodic Table is a weighted average of
the masses (in amu) of its isotopes found
on Earth.
Periodic Table
• First created by Mendeleev
• Still used today
• Started by arranging elements by weight
beginning with hydrogen
• Noticed trends so he started arranging them in
rows (periods) and started a new period every
time he found an element with properties like
hydrogen.
• Discovered that elements in other groups had
similar properties as well.
Definitions
• Main-group elements- elements in
groups 1A, 2A, and 3A-8A.
• Transition elements- elements in B
groups
• Inner transition elements- elements 5871 and 90-103
Classes of Elements
• There are three classes
• Metals- most elements in the Table. They
are solid (except Hg), shiny, conductors,
ductile (wires), and malleable (sheets)
• Alloys- solutions of one or more metals
dissolved in another metal. Examples:
Bronze (copper and tin) Pewter (tin,
antimony, and lead)
Classes (cont)
• Nonmetals- with the exception of
Hydrogen, they are on the right side of the
table.
• Non conductive (exception graphite)
• Tend to accept electrons in reactions
• Metalloids- B, Si, germanium, arsenic,
antimony, and tellurium
• Have some properties of both
Trends in the Periodic Table
• Sizes increases as you move down a
column
• For the halogens, group 7A, MP and BP
increase as you move top to bottom
• Ionization energy- is a measure of how
difficult it is to remove an electron from an
atom in the gaseous state
• Ionization energy increases as you move
up a group and from left to right across a
period
Electron Configuration
• Electrons are located in the considerably
larger space outside the nucleus.
• The lowest possible energy level an
electron can be in is the Ground State.
• The energy of electrons is quantized and
there are only certain energy levels an
electron can be in.
Electron Configuration
• Electrons in atoms are confined to specific
regions of space called Principal energy
levels or more simply, shells.
• Shells are numbered 1,2,3,4, and so on
• Electrons in the first shell are closest to
the nucleus, are held most strongly, and
therefore are the lowest in energy.
• As you move away, energy increases
Electron Configuration
• Shells are divided into subshells
designated by the letters s, p, d, and f.
• Within subshells electrons are grouped
into orbitals.
• Orbital- is a region of space that can hold
two electrons.
• The first shell has one orbital called the 1s
• The second shell contains one s orbital
and three p orbitals.
Electron Configuration
• P orbitals come in sets of three and holds
a total of six electrons
• The 3rd shell contains one s orbital, three
p orbitals, and five d orbitals.
Orbital Shapes
• All s orbitals have the shape of a sphere
with the nucleus in the middle.
• Each 2p orbital has the shape of a
dumbbell with the nucleus at the midpoint
of the dumbell.
• The three 2p orbitals are at right angles to
each other.
Electron Configuration
• Electron Configuration- a description of
the orbitals that its electrons occupy
• In the ground state of an atom only the
lowest energy orbitals are occupied, all
others are empty
• We determine the ground state
configuration using the following rules:
Rules
• Orbitals fill in the order of increasing
energy from lowest to highest
• Each orbital can hold up to two
electrons with spins paired
• When there is a set of orbitals of equal
energy, each orbital becomes half-filled
before any of them becomes
completely filled.
3 types of configuration notation
• Orbital box diagrams- use boxes or lines
to represent an orbital
• Condensed configuration- write only
shell number, orbital, and number of
electrons in that orbital
• Noble Gas Notation- use noble gas
symbol for inner shell designation
Showing electron config.
• Usually only show the outer shell electrons
• Outer-shell electrons are called valence
electrons
• The energy level in which valence
electrons are found is called the valence
shell.
• To show the outmost electrons of an atom,
we commonly use a representation called
a Lewis Dot Structure
Lewis Dot Structures
• Lewis Dot Structures- shows the symbol
of the element surrounded by a number of
dots equal to the number of electrons in
the outer shell of an atom of that element.
• Valence electrons- are electrons in the
outermost shell.
• Examples-
Another trend in the Periodic Table
• Elements in a group have the same
electron configuration in their outer shell!!
• That mean the Lewis Structure will be the
same as well, except for the atomic
symbol.
Atoms vs. Ions
• Atoms are neutral species that have the
same number of electrons and protons.
• During typical chemical reactions, the
number of protons and neutrons does not
change
• The number of electrons does change!!
Atoms can gain or lose electrons.
• Ions are a compound or element that has
gained or lost an electron and is now
charged!!
Atoms vs. Ions
• Example-
• Ionization Energy- a measure of how
difficult it is to remove an electron from an
atom.
• Ionization Energy increases as you move
up a group and from left to right on the
periodic table.