Download Atoms

The Study of Matter
What is Chemistry?
What are elements?
What are compounds?
What is the periodic table?
Atomic Theory?
Phosphorous Spontaneously reacts
with oxygen and bursts into
flames
A Piece of
Magnesium…..on fire.
Flame Test
• Chemistry: the physical science that deals with the composition,
properties, and changes in matter.
• Observation: a direct form of knowledge obtained by means of
one of your five senses—seeing, smelling, tasting, hearing, or
feeling.
• Interpretation: an indirect form of knowledge that builds on a
concept or an experience to further describe or explain an
observation.
 Observable knowledge is called empirical knowledge
 Theoretical knowledge explains and describes scientific
observations in terms of ideas; theoretical knowledge is not
observable.
Matter
Pure Substances
Mixtures
Mechanical mixtures
Elements
Compounds
Solutions
Suspensions
Colloids
Mixtures: Homogenous and Heterogeneous
 Mechanical mixtures have two separate identities that
can be distinguished visually
 ex. Salt mixed with sand
 Solutions a solute and a solvent combine to form a
solution that does not have any obvious
characteristics of the previous two.
 The solutions physical and chemical properties may vary a
great deal from previous properties of the solute and
solvent
Matter
Pure Substances
Mixtures
Mechanical mixtures
Elements
Compounds
Solutions
Suspensions
Colloids
Pure Substances
 Pure substances can be subdivided into two
groups also; compounds and elements
 Compounds a combination of two or more
elements.
Ex: Carbon dioxide (CO2)
chemical formula
 Elements contain a single kind of atom.
Ex: Sodium (Na)
 So now that we know about mixtures and pure
substances lets look at them on a smaller scale.
 What is it exactly that makes up these elements, that in
turn can make compounds, which in turn can make
mixtures?
• The atom is the smallest particle of an
element that maintains the characteristics of
the element
• Atoms are composed of a positive nuclei
surrounded by a cloud of negative particles
 Inside the nuclei: protons
(positively charged ions)
and neutrons (electrically
neutral ions)
 Together protons and
neutrons are referred to as
nucleons
 Outside the nuclei: electrons
(negatively charged ions)
that randomly circulate
around the nuclei
• The number of protons in the nucleus will determine
what element we have
• How can we organize and classify these different
elements?
John Alexander Newlands:
• Arranged elements in order of increasing atomic masses (we
will cover this shortly)
• Noticed some reoccurring properties and called this periodic
law
Dmitri Mendeleyev:
• Published the periodic table of elements
• Left spaces for elements not yet discovered
•
http://zircon.mcli.dist.maricopa.edu/mlx/warehouse/00601-00700/00696/Chemistry.swf
• Family/Group: vertical columns
• Have similar chemical properties
• Period: horizontal rows
• These elements change from metals to nonmetals
• The furthest left is the most metallic and the
furthest right is the least metallic
• Metals: usually solid at room temperature, good
conductors of heat and electricity, bendable, malleable
(can be pounded into sheets), shiny and ductile (can be
stretched)
• Non-metals: solid, liquid or gas at room temperature,
poor conductors of electricity, brittle to the touch and
lack the lustre of metals
Red
Phosphorous
• Metals
• Nonmetals
• Metalloids (semi metals)
• Main Group Elements
• Transition Metals
• Inner Transition Metals
 Alkali Metals:
 Group 1: IUPAC or (IA: American)
 Soft, silver coloured metals that react violently with water to
form basic solutions.
 Most reactive are cesium and francium
www.periodicvideos.com
Sodium demo
• Alkali Earths:
• Group 2 (IIA) www.periodicvideos.com
• Light, reactive metals that form oxide
coatings when exposed to air
 Transition Metals:
Groups 3-12 or IB-VIIIB (see inside cover of text)
Have more than one charge and exist in nature
as more than one species
 Metalloids:
 All the elements on the LEFT side around the
staircase.
• Halogens
• Group 17 or VIIA
• Highly Reactive
• Noble Gasses
• Group 18 or VIIIA
• Very Non-reactive
Inner Transition Metals
 Lanthanides
Atomic numbers 58-71
 Actinides
Atomic numbers 90-103
Dalton
Rutherford
Thompson
Neils Bohr
Law of Definite Proportions:
Elements combine in ratios
1:1
2:1
1:3
Lots of
Combinations
Law of Multiple Proportions
Elements with multiple charges for different
combinations
Law of Conservation of Mass
Matter: Atoms, Elements, molecules cant be created
or destroyed.
Thompson:
 Discovered the electron
 He proposed the atom was a positive
sphere with negative electrons pushed
into it like chocolate chips in a cookie.
***Raisin Bun Theory***
• Rutherford:
• His “gold foil” experiment led to the discovery of the proton
• Proposed that all atoms have a positive central region called the
nucleus
http://www.waowen.screaming.net/revision/nuclear/rsanim.htm
• Bohr:
• Electrons orbit the positive nucleus in certain orbits –
electrons cannot exist between orbits
• The higher the energy level of an electron, the further it is
from the nucleus
Steps For Drawing Bohr Diagrams:
1.
2.
3.
Write the Atomic symbol for the element and its name underneath it.
Write the number of protons above the atomic symbol…Atomic Number
Draw a line and counting up to the number of protons add electrons to each
level, separating each level with lines.
***Remember 2,8, 8, 18, 32***
http://youtube.com/watch?v=vUzTQWn-wfE
The atom song
• The number of protons in the nucleus of an atom.
• Is equal to the number of electrons in an atom.
• The basis for the order of elements on the periodic table.
© Addison-Wesley Publishing Company, Inc.
• mass # = protons + neutrons

always a whole number

NOT on the
Periodic Table!
© Addison-Wesley Publishing Company, Inc.
• Atoms of the same element with different mass numbers.

Nuclear symbol:
Mass #
Atomic #

Hyphen notation: carbon-12
12
6
C
© Addison-Wesley Publishing Company, Inc.
• Chlorine-37
•atomic #:
17
•mass #:
37
•# of protons:
17
•# of electrons: 17
•# of neutrons: 20
37
17
Cl
• Atomic Number:
• Refers to the number of protons an atom contains
• This number corresponds to the number of each
element on the periodic table.
• It’s the number of protons that determines which
element an atom is!
• Mass Number:
• Refers to number of protons plus neutrons in an
atom.
• This number closely resembles the atomic molar
mass found on the periodic table.
• In an atom the number of electrons and
protons are always equal.
• This makes an atom electrically neutral
• The Protons + Neutrons make up about 99%
of an atoms mass (electrons are
negligible…very very small)
• But, Electrons occupy most of the volume of
an atom.
What is the mass number of sodium?
How many electrons does an oxygen atom contain?
A carbon atom always contains 6 protons. If it has
an atomic mass of 13 how many neutrons are found
in the nucleus?
If a nitrogen atom has 8 neutrons in its nucleus what
is the atomic number?
Try It Out
Protons:
20
10
80
35
Ne
Br
10
Neutrons:
10
Electrons:
10
Protons:
35
Neutrons:
45
Electrons:
36
• Molecular Elements are composed of two or more
identical atoms ie. Bromine = Br2
• Each element usually has a particular electrical charge
associated with it.
• These charges indicate the number of electrons that
may be given up or accepted by a particular element
in order to form compounds.
• When atoms gain or lose electrons they form charged
Ions.
Ions are created when elements lose or gain electrons in
their valence orbital.
When we were drawing Bohr diagrams of elements the
last electron level that there wasn’t enough electrons to fill
was the Valence Orbital.
(8) X
Which is the Valence Orbital?
(8) √
(2) √
(8) X
(8) √
(2) √
(8) X
(8) √
(2) √
Negatively
Charged Ions
are called
Anions.
e
(8) √
(2) √
Positively
Charged Ions
are called
Cations.
(8) √
(8) √
(2) √
If you made Bohr diagrams for all the non metals you would
notice this.
(-3)(-2)(-1)
Write it in your data booklet!!!
To become stable the atoms will gain
or lose electrons (whichever is
easiest) to fill the outer energy level.
• When they lose electrons, they become
positively charged (cations)
• When they gain electrons, they
become negatively charged (anions)
• Polyatomic ions are single ions composed of multiple
atoms
Naming Ions:
If the ion is a non-metal, change the
ending to “ide”
If the ion is a metal, call it the same
name and add the word “ion” after it
******Polyatomic ions are as they
appear on your periodic table******
Calculating Charge
1. Find the noble gas that the element is closest to on the
periodic table (does not have to be in the same period)
2. Calculate if it would be easier to gain or lose electrons to get
the same number of electrons as the closest noble gas.
3. Calculate how many electrons the element would gain or
lose. (When you gain electron the ion becomes positive, when
you lose, the ion becomes negative)
4. Record the charge as a superscript following the element’s
symbol.
1)
Find the noble gas that the element is
closest to on the periodic table (does
not have to be in the same period)
2)
Calculate if it would be easier to gain
or lose electrons to get the same
number of electrons as the closest
noble gas.
3)
Calculate how many electrons the
element would gain or lose. (When you
gain electron the ion becomes positive,
when you lose, the ion becomes
negative)
4)
Record the charge as a superscript
following the element’s symbol.
O
Gain electrons
2 electrons
O2-
Calculating Charge
Oxygen
+1
+2
Practice
2-
O
Closes noble gas is
2 spots away.
3-
2-
1-
• Three classes of compounds are
possible.
• Ionic compound – a metal-nonmetal
combination (cation/anion)
• i.e. Na+ and Cl-  NaCl(s)
•Molecular compound – a nonmetalnonmetal combination
•i.e. H2 and O2  H2O(l)
•Inter-metallic compound – a metalmetal combination
•i.e. Brass  CuZn(s)
• A diagnostic test can determine if a
solution is ionic or molecular.
•Ionic - likely to conduct electricity
• This is because ionic cmpds exist
as ions in solution
NaCl(s)  Na+(aq) + Cl-(aq)
Conducts electron energy
•Molecular - unlikely to conduct
electricity
• This is because molecular cmpds
exist as molecules in solution
C6H12O6(s)  C6H12O6 (aq)
Ionic compounds are solid at room
temperature, while molecular
compounds can be solid, liquid or
gas at room temperature.
• Atoms form ions to become more
stable. (full energy levels)
• Metals want to lose electrons and
nonmetals want to gain electrons.
• Ionic bonds are formed when a
metal gives its Valence electrons to
a non-metal.
• Metals become positively charged
ions called cations.
• Nonmetals become negatively
charged ions called anions.
• The simplest type of ionic
compounds are called binary ionic
compounds. These compounds have a
metal and a nonmetal.
8
+
-
Sodium Chloride
(NaCl)
1. Write out the ions present in the compound.
Al3+ and O22. Determine the number of each ion to produce a
neutral compound. Find the coefficient of each
element that, when multiplied by the charge, will
provide the lowest common multiple.
So what do we have to
multiple 3 by to make it
3+ and
Al
6?
2
O23
So what do we have to
multiple 2 by to make it
6?
So using multiplication, when is the first time the 3 times table
and the 2 times table overlap?
6
3. Write out the subscripts all together.
Al2O3
***In Ionic compounds we are balancing charges
and are therefore recording the simplest wholenumber ratio of ions
The “criss cross” method can also
be used.
But you have to understand WHY it works!!!
Lithium & Oxygen
• Li2O
Barium & Phosphorus
• Ba3P2
Magnesium & Fluorine
• MgF2
Sodium and Chloride
• NaCl
• Metal (the first element in the
compound) uses its full name
• Nonmetal is names using the suffix
“ide”
• Prefixes are not necessary when
naming ionic compounds because
there is only one possibility
between a metal and nonmetal.
potassium oxide
magnesium chloride
actinium sulfide
K2O (s)
MgCl2 (s)
Ac2S3 (s)
• There are 2 types of ions that make this
process slightly more complex
Multi-Valent Metals
(More than 1 Charge)
Polyatomic Ions
• Some metals (transition metals) can
form more than 1 ion. (every ion has
a charge)
• Because of this, these elements
have more than one charge……listed
on the periodic table.
Iron
Fe
Fe
3+
2+
Cobalt
Co
Co
2+
3+
Chromium
Cr
Cr
3+
2+
Mercury
Hg 2+
Hg +
• The first charge listed in the periodic table is the most common.
If the charge is not provided assume we are using that charge.
• Iron can form two different
compounds with oxygen.
Fe2O3 & FeO
• When given a chemical formula use
the non-metal to determine the
appropriate charge
Example: Cu2O
Which charge of copper was used?
+2 or +1?
-2
Cu2
O
+1
= 2*+1
+2
+2
= 2*+2
+4
• When naming a compound that contains a
multivalent element be sure to indicate the charge
that was used by putting it in brackets after naming
the particular element
• Ex: Iron (III) Oxide vs. Iron (IV) Oxide
Number
1
2
3
4
5
6
7
Roman Numeral
I
II
III
IV
V
VI
VII
Step 1: Name the metal and place the
metallic ion charge (Roman numerals)
in parenthesis after the name.
Step 2: Name the nonmetal ion as
usual.
iron (II) oxide
Fe2O3(s)
nickel (II) sulfate
Ni2(SO4)3(s)
• Atoms can bond to form very
strong, stable polyatomic ions which
have a net positive or negative
charge.
• This net cluster of atoms act as a
unit.
• It participates in chemical reactions
without breaking apart.
• Many polyatomic ions end in –ate,
-ite, -ide, -oate.
i.e. sulphate ion = SO42hydroxide ion = OHFound as a separate table on your
periodic table or on the back cover
of your textbook.
• Metal receives its full name, and the
polyatomic nonmetal receives the
polyatomic name (found on the
periodic table)
• *please note that there is a
polyatomic ion named ammonium –
NH4+
NaClO3(s)
Mg(HSO3)2(s)
ammonium hydroxide*
• Brackets need to be used around
polyatomic ions when multiplied.
•Example: Ca2+ & NO3
Ca(NO3)2
Mg2+ &
SO42-
MgSO4
Ca2+ &
Ca(NO3)2
H+
&
BO33-
H3BO3
NO3-
Ba2+ & PO43Ba3(PO4)2
• Water is bonded (loosely) with an
ionic compound.
• When hydrates are heated, the
bond to the water is broken
releasing water to the atmosphere
as water vapour.
• When calculating the molar mass
include the water as well
CuSO4∙6H2O(l) + heat  CuSO4
Copper (II) Sulfate water (1/6)
(s)
+ H2O(g)
CuSO4∙6H2O(l)
Cu
S
+O
1 X 63.55 g/mol
1 X 32.07 g/mol
4 X 16.00 g/mol
159.62 g/mol
H
+O
2 X 1.01 g/mol
1 X 16.00 g/mol
6 X 18.02 g/mol
108.12 g/mol
267.74 g/mol
• 1. Use prefixes to indicate the number of water
molecules attached to the ionic compound
Copper (II) SulfatePenta Hydrate
CuSO4 ∙ 5H2O(s)
Sodium CarbonateDeca Hydrate Na2CO3 ∙ 10H2O(s)
2. Add a fraction at the end to indicate how many waters
attach to one molecule
CuSO4 ∙ 5H2O(s)
Na2CO3 ∙ 10H2O(s)
3. Add a fraction at the end to indicate how many waters
attach to one molecule
Copper (II) Sulfate 1/5 Hydrate
CuSO4 ∙ 5H2O(s)
Sodium Carbonate 1/10 Hydrate Na2CO3 ∙ 10H2O(s)
• Occurs when two nonmetals share
electrons
• Both nonmetals want to gain
electrons to attain a noble-gas
type structure.
• Neither atom wants to lose
electrons, so they compromise by
sharing the electrons.
O2
F2
Cl2
• To name molecular
compounds, one must
use prefixes to
express the number
of atoms of each
nonmetal.
mono
di
tri
tetra
penta
hexa
hepta
octa
nona
deca
1
2
3
4
5
6
7
8
9
10
 It is necessary to use prefixes for
nonmetals because there may be different
compound using the same two nonmetals.
 For example: nitrogen and oxygen can
combine in several different ways:
nitrogen
monoxide*
NO(g)
nitrogen dioxide
NO2(g)
nitrogen trioxide
NO3(g)
*Please note that the prefix “mono” is never
used on the first atom*
• Some molecular compounds have
“common names”
ammonia
NH3(g)
ethanol
C2H5OH(l)
glucose
C6H12O6(s)
hydrogen peroxide
H2O2(l)
hydrogen sulphide
H2S(g)
methane
CH4(g)
methanol
CH3OH(l)
octane
C8H18(g)
ozone
O3(g)
propane
C3H8(g)
sucrose
C12H22O11(s)
water
H2O(l) or HOH(l)
Acids
Bases
Taste sour
Taste bitter
Turn litmus paper red
(diagnostic test)
Turn litmus paper blue
(diagnostic test)
Have pH less than 7
Have pH more than 7
Neutralize bases
Neutralize acids
Composed of nonmetals
Composed of metal and
nonmetal
• Acids cannot be classified as ionic
or molecular because they have
characteristics of both.
• Acids are usually composed of all
nonmetals which would indicate it is
a molecular compound.
• Conversely, acids conduct
electricity which would indicate
that they are ionic compounds.
• Acids show their properties when
they are dissolved in water. They
are given the symbol (aq).
HI(aq)
HCl(aq)
HBr(aq)
H2SO4(aq)
HNO3 (aq)
HClO4(aq)
)
)
Hydro Bromic Acid
Hydro Iodic Acid
Sulfuric Acid
Perchloric Acid
Nitric Acid
Hydro Chloric Acid
HI(aq)
HCl(aq)
HBr(aq)
H2SO4(aq)
HNO3 (aq)
HClO4(aq)
)
)
Hydro Iodic Acid
Hydro Chloric Acid
Hydro Bromic Acid
Sulfuric Acid
Nitric Acid
Perchloric Acid
• Acids can be identified because the
hydrogen ion is written first in the
formula or ends in “COOH”.
• If the compound does not have the
aqueous symbol after the formula, it
is not named as an acid, but rather
an ionic compound. For example,
HCl(g) is named as hydrogen chloride.
• The most recent naming system for
acids is to name the acidic compound
as you would any other ionic compound,
but place the word “aqueous” in front
of the name.
Aqueous hydrogen sulfate
H2SO4(aq)
Aqueous hydrogen chloride
HCl(aq)
• The classical rules for naming acids are often used
today. There are 3 classical rules:
hydrogen ______ide
Chlor ⃗ hydro____ic
Chlor acid HCl
hydrogen ______ate
⃗ ________ic
Chlor
Chlor acid HClO3
hydrogen ______ite
acid HClO
Chlor
Chlor ⃗ _______ous
2
sulfuric acid
hydrochloric acid
chloric acid
chlorous acid
H2SO4(aq)
HCl(aq)
HClO3 (aq)
HClO2 (aq)
• The strongest bases are soluble ionic hydroxides
• ie: sodium hydroxide, barium hydroxide
NaOH
Ba(OH)2
Bases have a pH of greater than 7. (7-14).
Strong bases that contain OH have a pH of around 11-14