Download U1 Atoms, Elements and Ions

Sec. 0.3: Chemical Foundations
Elements, Atoms,
and Ions
3-1
Objectives
• Learn the names and symbols for some
elements.
• Learn about the relative abundance for
some elements.
• Learn about Dalton’s theory of atoms.
• Understand the law of constant
composition.
• Learn about how a formula describes a
compound’s composition.
• Understand Rutherford’s experiment and
its impact on atomic structure.
3-2
Objectives
• Describe important features of subatomic
particles.
• Learn about isotope, atomic number and
mass number.
• Understand the use of the symbol X.
• Learn the various features of the periodic
table.
• Learn the properties of metals, nonmetals,
and metalloids.
• Describe the formation of ions from their
parent atoms.
3-3
Objectives
• Predict which ion a given element forms by
using the periodic table.
• Describe how ions combine to form neutral
compounds.
3-4
Section 3.1: The Elements
• Remember, elements are combined to form
molecules the way letters are combined to
form words.
• Presently there are about 115 known
elements.
• Only 88 occur naturally, the rest are
made in laboratories.
• Only 9 elements account for most of the
compounds found in the Earth’s crust.
3-5
The Elements
• Scientists use the word element in many
different ways.
• Sometimes it is referred to in the
microscopic sense:
– A single atom of Au or Ag could be referred
to as an element.
– Also, molecules such as O2 or N2, are
referred to as elements.
• In the macroscopic sense we can refer to
a bar of “pure” iron or a 24k gold ring as
elements.
3-6
The Elements
• When we say something contains a
particular element we do not necessarily
mean free atoms, but may also mean in a
form combined with other elements in
some compound.
• Our bodies contain many “trace” elements
– elements that are present in very small
amounts, but are crucial to life.
• Some of these elements include: arsenic,
chromium, cobalt, copper, fluorine, iodine,
manganese, molybdenum, nickel, selenium,
silicon and vanadium.
3-7
Section 3.2: Symbols For The
Elements
• Just as each state has a two-letter
abbreviation, each element has a one- or
two-letter symbol to make life simple for
chemists.
• The list of trace elements from the previous
slide can be simplified to: As, Cr, Co, Cu, F,
I, Mn, Mo, Ni, Se, Si, & V.
• Notice the first letter is ALWAYS
capitalized and the second letter, if
present, is NEVER capitalized.
3-8
Symbols For The Elements
• Some symbols make sense like O for
oxygen and H for hydrogen or Ni for
nickel.
• Others, like Pb for lead or Fe for iron,
don’t automatically make sense; they
originated from the Greek or Latin names
of plumbum (Pb) and ferrum (Fe).
• The only real way to learn them all is to
memorize them. Chemists have arranged
them in a periodic table to help with that.
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While there
are well over
100 different
elements,
many are
fairly rare; we
should know
the most
common
elements.
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Section 3.3: Dalton’s Atomic
Theory
• Scientists studying matter in the
eighteenth century made the following
observations:
– Most natural materials are mixtures of pure
substances.
– Pure substances are either elements or
combinations of elements called compounds.
– A given compound always contains the same
proportions (by mass) of the elements.
3-11
Dalton’s Atomic Theory
• One particular English scientist named
John Dalton attempted to explain these
observations in 1808.
• Dalton made his living as a teacher in
Manchester, England; he started a school
in his town when he was just 12 years old!
• He never married, was colorblind (to red)
and liked to bowl every Thursday
afternoon.
3-12
Dalton’s Atomic Theory
1. Elements are made of tiny particles called atoms.
2. All atoms of a given element are identical.
3. The atoms of a given element are different from
those of any other element.
4. Atoms of one element can combine with atoms of
other elements to form compounds. A given
compound always has the same relative numbers
and types of atoms.
5. Atoms are indivisible in chemical processes. That
is, atoms are not created or destroyed in chemical
reactions. A chemical reaction simply changes the
3-13
way atoms are grouped together.
Section 3.4: Formulas of
Compounds
• The types of atoms and the number of
each type in each unit (molecule) of a
given compound are conveniently expressed
by a chemical formula.
• The atoms are indicated by their symbols
and the number of each type is indicated
by a subscript (unless there is only one).
– Ex) C6H12O6 or H3PO4
3-14
Practice
•
Write the formula for each of the
following compounds, listing the elements
in the order given:
a. A molecule contains four phosphorous atoms
and ten oxygen atoms.
b. A molecule contains one uranium atom and six
fluorine atoms.
c. A molecule contains one aluminum atom and
three chlorine atoms.
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Section 3.5: The Structure of
the Atom
• In the late 1890’s J.J. Thomson, an
English physicist, determined atoms of any
element could be made to emit tiny
negative particles.
• He showed they were repelled by the
negative part of an electric field.
• He had discovered electrons.
• He also concluded atoms must contain
positive charges to cancel out the negative
(atoms are electrically neutral).
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The Structure of the Atom
• Building on J.J. Thomson’s discoveries,
Lord Kelvin proposed a “plum pudding”
model.
• Consider pudding with raisins in it. Now,
imagine the raisins are negatively-charged
and the pudding is positively-charged.
• The positive charge was cancelled out by
the negative charges giving an overall
charge of zero.
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Figure 3.3: Plum Pudding model of an atom.
3-18
The Structure of the Atom
• In 1911 another physicist named Ernest
Rutherford performed his famous “Gold
Foil Experiment” that concluded the plum
pudding model could not be correct.
• His experiment involved firing alpha
particles at a sheet of thin gold foil.
• The foil was surrounded by a detector
coated with a substance that produced
tiny flashes when hit by an alpha particle.
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Figure 3.5: Rutherford’s experiment.
3-20
The Structure of the Atom
• Since alpha particles are about 7500 times
more massive than electrons, Rutherford
expected them to tear through the foil
the way a bullet would go through paper.
• What actually happened was most passed
straight through, but some were deflected
at great angles and even reflected
backwards!
• According to the plum pudding model he
expected ALL to pass through with a few
being deflected VERY slightly.
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Figure 3.6: Results of
foil experiment if
Plum Pudding model
had been correct.
3-22
The Structure of the Atom
• He concluded the positively-charged alpha
particles could have only been deflected
by a dense center of concentrated positive
charge.
• He came up with the model of a nuclear
atom: a nucleus composed of positivelycharged protons orbited by negativelycharged electron made up of mostly empty
space.
• He eventually came up with idea of
neutrons also residing in the nucleus.
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Figure 3.6: Actual results.
3-24
The Structure of the Atom
• Ultimately, Rutherford’s model gave us the
three subatomic particles: protons (charge
= +1) and neutrons (charge = 0) in the
small central nucleus and tiny electrons
(charge = -1) orbiting the nucleus.
• If an atom were blown up to the size of a
professional football stadium, the nucleus
would be the size of a fly on the 50-yard
line.
• If the nucleus were the size of a grape,
the atom’s radius would be about a mile.
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The Structure of the Atom
• J.J. Thomson discovered electrons using a
cathode ray tube (CRT).
• CRT’s are sealed glass tubes containing a
gas and a metal plate at each end
connected to external wires.
• When an electric current was applied to
the plates a glowing beam was produced;
he was convinced it was a stream of
electrons.
• This technology is still used today in
televisions and computer monitors.
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Figure 3.7: Schematic of a cathode ray tube.
3-27
Section 3.6: The Modern
Concept of Atomic Structure
• Today, the view of the atom is:
– A tiny nucleus about 10-13 cm in diameter.
– Electrons that move around the nucleus at an
average distance of about 10-8 cm away.
– Electrons and protons having equal and
opposite charges while neutrons have no
charge.
– Protons and neutrons almost 2000 times more
massive than electrons.
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Figure 3.9:
A nuclear atom
viewed in cross
section.
3-29
Modern Atomic Structure
• Every atom is composed of the three basic
subatomic particles.
• Different elements have different numbers
of each of these particles.
• The reason one element behaves
differently than another lies in the number
and arrangement of their electrons.
• When atoms get close to each other their
electron “clouds” can overlap and interact.
• We’ll learn more about this later.
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Section 3.7: Isotopes
• We now know each element has a unique
number of protons and electrons (they
must be equal), but what about the
number of neutrons?
• Dalton assumed any two atoms of a given
element were identical; not quite correct.
• We can have two atoms of the same
element (same number of protons) with
different numbers of neutrons.
• These are called isotopes.
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Figure 3.10: Two isotopes of sodium.
3-32
Isotopes
•
There are two important numbers
associated with any given element:
1. Atomic Number – The number of protons in a
nucleus.
2. Mass Number – The SUM of the number of
protons AND neutrons (a.k.a. nucleons) in a
nucleus (NOT the sum of their masses).
•
We should note that two different
isotopes will have the same atomic
number, but different mass numbers.
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Isotopes
• Scientists like to use symbols as shorthand for
these terms:
– X = the symbol of the element
– A = the mass number (nucleons)
– Z = the atomic number (protons)
• A generic representation of any given element
would look as follows:
A
Z
X
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Isotopes
• The two previous examples of isotopes
of sodium would be:
23
11
Na
24
11
Na
•The example on the left would contain 11
protons and 12 neutrons (23-11=12).
•The example on the right would contain 11
protons and 13 neutrons (24-11=13).
3-35
Practice Problems
•
Write the symbol for each of the
following atoms, and list the number of
protons, neutrons, and electrons for
each.
1)
2)
3)
4)
The
The
The
The
cesium atom with a mass number of 132.
iron atom with a mass number of 56.
krypton atom that has 48 neutrons.
nitrogen atom that has 6 neutrons.
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Figure 3.11: The periodic table
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Figure 3.12: Elements classified
as metals and nonmetals.
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Figure 3.13: A collection of argon atoms.
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Figure 3.14: Nitrogen gas contains NXN molecules.
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Figure 3.14: Oxygen gas contains OXO molecules.
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Table 3.5
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Figure 3.15: The decomposition
of two water molecules.
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Figure 3.17: Spherical atoms packed closely together.
3-44
Figure 3.19: The ions formed by selected members of
groups 1, 2, 3, 6, and 7.
3-45
Section 3.8: Introduction to the
Periodic Table
Objectives:
To learn about various features of the
periodic table.
To learn some of the properties of
metals, nonmetals, and metalloids.
A Simple Version of the Periodic Table
3-47
• In any box on the Periodic Table, what
information can you find?
6
C
12.01
Average Atomic Mass = the
weighted average of all the
mass numbers for each
isotope of the element
Atomic number = number of
protons, unique for every
element, no 2 elements have
the same atomic #
Element symbol = can be
1,2 or 3 letters, first letter is
always capitalized, and
succeeding letters are always
lower case
3-48
Weighted Average Atomic Mass
•
•
Remember elements can have different isotopes which
means that they vary in their number of neutrons.
If you have 3 different isotopes of the same element:
– 15 atoms have a mass of 21
– 8 atoms have a mass of 23
– 2 atoms have a mass of 19
We can calculate the weighted average by multiplying the
number of atoms by their mass:
(15) (21) = 315
537 = 21.48
(8) (23) = 184
25
(2) (19) =+ 38
average atomic mass
537
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• If you have 10,000 atoms of Cl
– 7577 atoms have a mass of 35
– 2423 atoms have a mass of 37
What is the average atomic mass of Cl?
• (7577) (35) = 265195
(2423) (37) = 89651
354528
354528 = 35.45
10,000
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Using % to find Average Atomic Mass
• Usually we only know the percents of various
isotopes that make up different elements, we
can use this to calculate the average atomic
mass.
• If we have 100% chlorine:
75.77% of mass is 34.969 -> .7577x34.969 =
26.469
24.23% of mass is 36.966 -> .2423x36.966 =
8.95686
Add the 2 together to get the atomic mass:
26.469 + 8.95686 = 35.45
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Practice
• Oxygen has 3 isotopes 16O, 17O, 18O
99.76% of mass is 16O
0.04% of mass is 17O
0.20% of mass is 18O
What is the average atomic mass?
• Find the atomic mass if 99.64% of mass is 14N and
0.36% is 15N.
• Magnesium has three isotopes. 78.99% magnesium 24
with a mass of 23.9850 amu, 10.00% magnesium 25
with a mass of 24.9858 amu, and the rest magnesium
25 with a mass of 25.9826 amu. What is the atomic
mass of magnesium?
3-52
Periodic Table
• When looking at periodic table elements are
arranged in horizontal rows by increasing
atomic number.
• Horizontal rows are called “Periods”
Periods go left to right
As you move across the period the number of
valence electrons increases
3-53
Periodic Table
• The vertical columns are called “Groups” or
“Families”
• Elements in families share similar properties
• Each shares the same number of valence electrons in
outermost shell
• Can determine the number of valence electrons by the
number of the group
• Can use the group number and valence electrons to
find the charges of many elements
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Metals, Semimetals, Non-Metals
All elements on the periodic table are
grouped as metals, semimetals or
metalloids, or non-metals. Due to the
arrangement of the periodic table, it is
easy to identify each type of element.
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Metals: Fall to left and under the stairs
Properties of Metals:
Efficient conduction of heat and electricity
Malleability
Ductility
A lustrous appearance
Positively charged ions
3-56
Non-Metals: Right and above the stairs
• Dull, Brittle
• Negatively charged ions
• Nonconductors
-insulators
3-57
Semimetals or Metalloids: Makeup the stairs
• Properties of both metals and nonmetals
• Semiconductors
3-58
Lanthanide and Actinide Series
• Mostly human made elements
• Radioactive elements
3-59
Group 1A – Alkali Metals
•
•
•
•
Members of the group 1A
Have a +1 charge as ions
Very reactive with water (stored in oil)
Extremely malleable
and soft
3-60
Group 1A – Alkali Metals
• All members react with water to produce same ratio
of products:
Li + HOH -> LiOH + H2(g)
Na + HOH -> NaOH + H2(g)
K + HOH -> KOH + H2(g)
Rb + HOH -> RbOH + H2(g)
Cs + HOH -> CsOH + H2(g)
• As you move down the group:
– Reactivity with water increases
– Melting points and boiling points decrease
– Density increases
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Group 2A – Alkaline Earth Metals
• Members of Group 2A
• Form +2 charges as ions
• React with water
3-62
Group 2A – Alkaline Earth Metals
• The members of this group also react to form the
same ratio:
Mg + HOH -> Mg(OH)2 + H2
Ca + HOH -> Ca(OH)2 + H2
Sr + HOH -> Sr(OH)2 + H2
• As you move down the family we see some patterns
once again:
–
–
–
–
Reactivity increases
Melting/Boiling point decreases
Density increases
Slower to react than alkali metals…don’t have to store
them in oil
3-63
Trends for Group 1A & 2A
+1
+2
• Melting/Boiling points increase
• Reactivity increases
• Density increases
3-64
Group 3A or 13
•
•
•
Form +3 charge as ions
React with water in a 1metal:3 OH ratio
Slow reaction with water
3-65
Group 3A or 13
• When looking at Periodic Tables you will see this
group numbered 2 different ways, they are still
the same group that carries a charge of +3.
• Aluminum is a member of this family. It is a
metal, it is reactive. We don’t find pure
aluminum in nature, it is always coated with
aluminum oxide.
• Aluminum foil is coated with thin layer of
vegetable oil to keep shiny
3-66
Groups 4A,14 and 5A,15
• These groups have metals, nonmetals, and
metalloids which makes their properties more difficult to group.
• Group 4A can have a +4
or -4 charge
• Group 5A can have a -3 charge
3-67
Group 6A or 16: The Oxygen Family
• Form -2 charge as ions
• Members below oxygen form oxides with
putrid odor:
SO2, SeO2, & TeO2
3-68
Group 7A or 17: Halogen Family
• Exist as diatomic molecules
• Form ions with a -1 charge
• Non-metals
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Group 7A or 17: Halogens or Hydrogen Family
• Sometimes periodic tables will show
hydrogen in this family because it’s
property match better than the Alkali
metals.
• Members of this family are diatomic, they
are found as 2 bonded atoms: H2, F2, Cl2,
Br2, I2
• Members of this family are toxic and
poisonous (exception is H2)
• Ions of this family have -1 charge and are
called halides
3-70
Group 7A or 17: Halogens or Hydrogen Family
• F2 is toxic, and the most reactive element on the periodic
table, it can explode almost instantly with anything it mixes
with
• Cl2 is a poisonous gas, known as mustard gas in WWII
• Br2 is a red liquid with low vapor pressure, so it has a low
boiling point
• I2 is a silvery solid that sublimates from solid to gas.
Creates a purple toxic gas
• At2 radioactive and very rare, less than 30 grams on Earth,
not studied very often.
3-71
Group 7A or 17: Halogens or Hydrogen Family
• Halogens do not react with water, but we
can look at their reaction with an alkali
metal to find the pattern of reactivity for
the family.
– F2 + 2Na -> 2NaF explodes
– Cl2 + 2Na -> 2NaCl
vigorous
– Br2 + 2Na -> 2NaBr
slower
– I2 + 2Na -> 2NaI have to heat to get
reaction going
So what pattern is in the reactivity as you
move down the group?
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Group 7A or 17: Halogens or Hydrogen Family
As you move down the group:
• Boiling/Melting point increases
• Reactivity decreases
• Density increases
3-73
Group 8A or 18: Noble Gases
• Do not react easily with anything, due stable electron
configuration
• All other elements strive to reach noble gas
configuration for maximum
stability by reacting with
other elements.
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Group 8A or 18: Noble Gases
As you move down the period:
• Reactivity decreases
• Melting point increases
• Boiling point increases
3-75
Transition Metals: Group 3B-12B
• Have many electrons that they are able to share, this allows them
to have many different ions: Cr2+, Cr3+, Cr5+, Cr6+
• Always lose electrons to make positive ions
• These varying charges gives variety of
colors to same element
• Malleable and ductile
• Conduct electricity and Shiny
3-76
Charges of all the Families
• Remember atoms that lose or gain electrons form
ions, and these are the charges each family forms.
0
+1
+2
+3 +4 -3 -2
-1
Transition Metals
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A Few More General Periodic Trends
• As you move down a family atom size increases
• As you move left to right across the table
atom size decreases
• Where are the largest
atoms located?
3-78
3.9 Natural States of the Elements
Objectives:
To learn the natures of some common
elements.
Who is a solid, liquid or gas?
• When we look at the elements on the periodic
table, who is a solid, liquid or gas in their natural
state?
• Most elements are not found in their elemental
state, most elements are found in compounds with
other elements.
• Most elements on the periodic table are solids, so
we will point out those who are gas or liquid?
3-80
Liquids
• Only 2 elements in their elemental form are a
liquid at 25 degrees Celsius: Mercury and
Bromine
• Gallium and Cesium almost qualify, but they
are solids 25 degrees Celsius, but melt around
30 degrees Celsius.
3-81
Gases
• More elements exist in
their elemental form as a
gas, but there are some
important distinctions to
make about these gases.
• The noble gases are a gas,
called monatomic gas.
This means that the prefix
mono- means one. And
monatomic gases exist as
individual atoms.
Figure 3.13: A collection of argon
atoms.
3-82
Gases
• There is another group
of gases called diatomic
gases. The prefix dimeans two. These
elements travel in pairs
as molecules.
Figure 3.14: Nitrogen gas
contains NXN molecules.
Figure 3.14: Oxygen gas
contains OXO molecules.
3-83
Gases
There are 7 elements that exist as diatomic molecules, you
will simply need to find a way to memorize these.
If you notice, all of the halogens fall in this category, and
then hydrogen, nitrogen, and oxygen.
You will also notice that 2 of these are not gases, make sure
you do not for get to include these in your diatomic list.
3-84
Noble Metals
• There is one group of metals which are
relatively unreactive and thus called the noble
metal. This group includes gold, silver and
platinum.
3-85
Periodic Table
Diatomic Solids
Diatomic Liquids
Diatomic Gases
Monatomic Gases
Liquids
Solids
Synthetically Made
3-86
3.10 Ions
Objectives:
To describe the formation of ions from
their parent atoms and learn to name
them.
To predict which ion a given element forms
by using the periodic table.
What is an ion?
• When we discussed atoms before, we were
always looking at a neutral atom.
Neutral atoms always have equal numbers of
protons and electrons.
protons = +1 charge
electrons = -1 charge
• When atoms have unequal numbers of
protons and electrons, then the atom is a
charged particle called an ion.
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Ion Facts
• Ions are atoms, or groups of atoms, with a
charge.
• The charge is created by different numbers of
protons and electrons.
• In an ion ONLY electrons can move.
• Atoms gain or lose electrons to become ions.
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Cations and Anions
• There are 2 types of ions: cations and anions.
• Cations are ions with a positive charge. To form
a cation, an atom has lost electrons.
Example: Na loses an electron and becomes Na+
• Anions are ions with a negative charge. To form
an anion, an atom has gained electrons.
Example: Cl gains an electron and becomes Cl-
3-90
Basic Names for Ions
• Cations do not change names from their neutral
atoms.
Example: Magnesium loses 2 electrons and
becomes Mg2+ which is named magnesium ion.
• Anions change the end of their name to –ide.
Example: Chlorine gains an electron and
becomes Cl-. We would change the name from
chlorine to chloride.
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Some Common Anion Names
• What would the names of the following ions
be?
• Chlorine =
• Fluorine =
• Bromine =
• Iodine =
• Oxygen =
• Sulfur =
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How to Determine the Charge
• When determining the charge for an atom we
can use the periodic table to help.
• The number of valence electrons determines the
charge.
• All atoms want 8 valence electrons.
• If an atom has 1-3 valence electrons the atom
will lose them to become positive.
• If an atom has 6-8 valence electrons the atom
will gain electrons to become negative.
• We can determine the charge by looking at the
periodic table.
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Figure 3.19: The ions formed by selected members
of groups 1, 2, 3, 6, and 7.
3-94
Practice
• Determine the name and charge of the
following ions:
Potassium
Bromine
Calcium
Sulfur
Aluminum
Strontium
Cesium
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3.11 Compounds that Contain Ions
Objective:
To describe how ions combine to form
neutral compounds.
Ions and Compounds
• Any time a bond is formed between 2 or
more ions, we call this an ionic compound.
• In an ionic compound the overall charge must
be zero.
• This means there must be cations and anions
present.
Example: Na+ + Cl- -> NaCl
Each atom has a charge of +1 or -1, when we
add this together it comes out to be zero.
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• If we add two ions together and they are not equal,
then we must add another ion to balance the
charge.
Example: Mg2+ + 2Cl- -> MgCl2
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