Atoms and Bonding
Chapter 4
Unit – Introduction to the chemistry
Physical Science
Mrs. N. Castro
Homework
 Work chapter’s vocabulary words in your home.
 Remember, as always, do it in the index cards.
 Front part – the vocabulary word.
 Back part – vocabulary word definition.
 Bring all the index card with a rubber band for tomorrow.
 Value – 10 points.
Atoms, bonding and the Periodic Table
 Read pages 125 to 127.
 Answer the following:
 What is compound?
 How the elements are
combined?
 What is chemical bond?
 Which particle from the
atom can form chemical
bond?
 Compound – Substance made
of two or more elements
chemically combined in a
specific ratio or proportion.
 The elements are combined
using a chemical bond.
 Chemical bond – force of
attraction that holds atoms
together as a result of the
rearrangement of the
electrons between them.
 Electrons
Atoms, bonding and the Periodic Table
 Is there all electrons in the
same energy levels?
 Which electrons can form
chemical bond, all of them or
any specific?
 In which energy level we can
find the valence electrons?
 What determines the valence
electrons in each atoms?
 No
 Valence electrons
 In the higher energy level
with the higher amount of
energy.
 They determines the
chemical property, to whom
the atom will establish a
chemical bond and how will
be.
Atoms, bonding and the Periodic Table
Group or Family number
Valence electrons
Chemical bond
1
1
1
2
2
2
13
3
3
14
4
4
15
5
3
16
6
2
17
7
1
18
8
0
Which is the maximum of chemical bonds that the atom can do?
Which family is able to perform the maximum chemical bond?
Which is the minimum?
Note: Families 3 to 12 will be study in future grades.
Atoms, bonding and the Periodic Table
 How the valence electrons can help me to determines the
quantity of chemical bonds the atom can perform?
 Using the electron dot diagram or Lewis diagram.
 To perform electron dot diagram:
1.
2.
3.
4.
5.
6.
Write the element symbol.
Surround the symbol by dots.
Each dot mean one valence electron.
Maximum numbers of dots – 8
Maximum dots per side – 2
Side - 4
Atoms, bonding and the Periodic Table
 Atoms tend to be more stable if they have 8 valence electrons.
 The only column in the Periodic Table with 8 valence electrons is
family # 18. They are the most stable elements in the Periodic
Table. This is the reason to explain why they don’t perform
chemical bond with any element.
 Family # 18 is nonreactive or stable.
 He (Helium) is an exception, because it has only 2 valence
electrons and it stable or non reactive element. It is part of the
Family 18.
 Another exception is H (Hydrogen), need only two valence
electron to be stable. Actually, it has only one valence electron.
Need only one more to complete it stability or nonreactive state.
Atoms, bonding and the Periodic Table
Atoms, bonding and the Periodic Table
 Ways to forms atoms bonds:
Valence electrons may be transferred from one atom to
another. One atom loose the valence electron and the other
gain it.
2. Atoms share valence electrons.
1.
Depend on how the atom form the bond, it will be the
chemical bond.
 Types of chemical bonds;

1.
2.
3.
Ionic bonds – metal with nonmetal
Covalent bonds – nonmetal with nonmetal
Metallic bonds – metal with metal
Atoms, bonding and the Periodic Table
 Metal
 React by loosing valence electrons.
 Reactivity will depend on how easily its atoms loose valence
electrons.
 Nonmetals
 React when they gain or share enough electrons to have 8
valence electrons (except H).
 Nonmetals usually combine with metals by gaining electrons.
 Nonmetals can also combine with other nonmetals and
metalloids by sharing electrons.
Atoms, bonding and the Periodic Table
 Metalloids
 Can either lose or share electrons when they combine with
other elements.
 Hydrogen
 Its has 1 valence electron – Family 1
 Is nonmetal.
 Properties very different from alkali metals.
 Share its electron when form compound with other nonmetals
to obtain a stable arrangement of 2 electrons.
Atoms, bonding and the Periodic Table
N
1.
2.
3.
4.
5.
6.
Element name?
Find the protons, electrons and neutrons.
Is N reactive or stable?
Mention two elements with properties similar to N. Explain
why they are similar.
How many bonds N can performed?
With which elements, N will prefer to establish chemical bond?
Why?
Atoms, bonding and the Periodic Table
Ionic Bonds
Lesson 2
Pages 131 to 137
Ionic Bonds
 To perform ionic bonds you need ions
 ION is an atom or group of atoms that has an electric charge.
Example: Cl-, Na+
 Types of ions:
 Negative ion – when atom gain an electron.
 Positive ion – when atom lose an electron.
Examples:
K = 1 valence electron, Family 1, lose the electron = K+
Br = 7 valence electrons, Family 17, gain 1 = Br _
Metals atoms are likely to lose electrons.
Nonmetals atoms likely to gain electrons.
Ionic Bonds
 Common Ions
 Monoatomic – ion made of one element
 Example: K+
 Polyatomic – ions made of more than one elements. Can be
positive or negative charges.
 Example: HCO3
- = one Hydrogen, one Carbon and three Oxygen =
one negative charge
 See table on page 132.
Ionic Bonds
 Ionic bond – is the attraction between two oppositely
charged ions.
 Ionic compound – resulting compound from ionic bond. The
ionic compound always is neutral, charge = 0. The total
positive charge of all positive ions equals to total negative
charge of all the negative ions.
Ionic Bonds Video
Ionic Bonds – Formulas
 What is the different between this?
K , KBr , K+ , BrK = element symbol
K+ = positive ion
KBr = chemical formula
Br- = negative ion
 Chemical formula - is a group of symbols that show the ratio of
elements in a compound.
KBr – ratio = one K and one Br
MgCl2 – ratio = one Mg and two Cl
 The formula tell you:
1. The elements that component the compound.
2. The ratio of each element in the compound (subscripts).
Ionic Bonds – Formulas
 Rules:
To write the formula for compound, always write the positive
charge first and then the negative charge.
2. Add subscripts that are needed to balance the charge. ONLY
SUBSCRIPTS.
3. If no subscripts is written, its is understood that the subscript
is 1.
1.
Example: NaCl = no subscripts for both elements; means: one Na and
one Cl.
4.
Formulas for compounds off polyatomic ions are written in a
similar way.
Ionic Bonds – Names
 Rules for ionic compounds:
1.
The name of the positive charge comes first, followed by the name
of the negative ion.
NaCl – Sodium chloride
MgBr – Magnesium bromide
2.
3.
4.
5.
The name of the positive ion is usually the name of a metal.
If the positive charge is a polyatomic ion, use the name of the
polyatomic ion. NH4Br – Ammonium chloride
If the negative ion is a single element, the end of its name changes
to –ide. MgO – Magnesium oxide
If the negative ion is polyatomic, its name usually ends in –ate or –
ite.
NH4NO3 – Ammonium nitrate.
NH4HCO3 –Ammonium bicarbonate
Ionic Bonds
 Exercises:
Teacher exercises
1. Find the ratio of each element in the compound.
2. Write chemical formula for ionic's compounds.
3. Name ionic's compounds.
Complete exercises from pages 131 to 135.
Ionic Bonds
 Properties of ionic compounds:
1.
2.
3.
4.
5.
6.
7.
8.
Hard.
Brittle crystals.
Forms solids by building up repeating patterns of ions. This three
dimensional arrangement is called crystals.
High melting points
When dissolved in water or melted, they conduct electric current.
Attraction between ions is very strong, so it takes a lot of energy to
break the bond.
When ionic compounds dissolve in water, they become ions again.
The charged particles carry electric current where the neutral ionic
crystals do not.
Ionic Bonds
 Extra exercises:
A fluorine (F) ion has a charge of 1-. An aluminum (Al) ion
has a charge of 3+.
I.
1.
2.
3.
4.
Draw for both atom the electron dot diagram.
Explain how the F and Al would exchange valence electrons to form an
ionic compound. Draw the explanation first and then explain in your
own words.
Write the compound’s chemical formula.
Name the compound.
II. A potassium ion has a charge of 1+. A sulfide ion has a charge
of 2-. What is the chemical formula for ionic compound and
the name.
III. Name the following compound MgO
IV. Perform exercises on pages 136 and 137.
Covalent Bonds
Lesson 3
Pages 138 to 145
Covalent Bonds
 It formed when 2 atoms share electrons.
 Usually are between nonmetals atoms.
 The attractions between the shared electrons and the protons
in the nucleus of each atom hold the atom together in a
covalent bond.
 The product of a covalent bond is called a molecule.
 A molecule is a neutral group of atoms joined by covalent
bonds.
Covalent Bonds
Covalent Bonds
 Exercise:
 Draw the Lewis diagram for Bromine atom.
 Draw the Lewis diagram for the Br2 molecule.
 Perform exercise on page 139.
Covalent Bonds
 Atoms can form single, double and triple covalent bonds by
sharing one or more pairs of electrons.
 Double bond – 2 atoms share 2 pairs of electrons
Covalent Bonds
 Triple bond – share 3 pairs of electrons.
Covalent Bonds
 Identify the single, double and triple covalent bonds:
Draw the Lewis diagram for (CH3)2SO - dimethyl sulfoxide
Covalent Bonds
 Solution
C
Fam. 14
H
S
O
Fam.1
Fam.16
Fam.16
Covalent Bonds
 Properties:
Do not conduct electric current when melted or dissolved in
water.
2. Lower melting point and boiling point.
1.
Covalent Bonds
Sometimes, atoms of some elements pull more strongly on
the shared electrons of a covalent bond than do atoms of
other elements. As a result, the electrons ser shared
unequally.
 Unequal sharing of electrons causes covalently bonded
atoms to have slight electric charges.

Covalent Bonds
 Types of covalent bonds:
1.
Non polar bond – a covalent bond in which electrons are
shared equally.
Example: H2
Covalent Bonds
2.
Polar bonds – Covalent bonds in which electrons are share
unequally. When electrons in a covalent bond are shared
unequally, the atom with the stronger pull gains a slightly
negative charge. The atom with the weaker pull gains a
slightly positive charge.
Example: HF
Covalent Bonds video
Bonding in Metals
Lesson 4
Pages 146 -151
Bonding in metals
 Metal atoms lose electrons easily because they do not hold





their valence electrons very strongly.
The loosely held valence electrons in metal atoms result in a
type of bonding that happens in metals.
Most metals are crystalline solids.
A metal crystal is composed of closely packed, positively
charged ions.
Each metal ion is held in the crystal by a metallic bonds.
Metallic bonds is an attraction between a positive metal ion
and electrons surrounding it.
Bonding in metals
 Metallic bonds is formed an attraction within metal atoms.
 Video ...\..\Clark-9805.Fig.4 (2).mov
 Properties:
Luster – When the light strikes these valence electrons, they
absorb the light and then re-emit the light.
2. Malleability and Ductility – The positive metal ions are
attracted to the loose electrons all around them rather than to
other metal ions. These ions can be made to change position.
However, the metallic bonds between the ion and the
surrounding electrons keep the metal ions from breaking
apart from one another.
1.
Bonding in metals
3.
4.
Thermal conductivity – Metals conduct heat easily because the
valence electrons within a metal are free to move.
Electrical conductivity – Metals conduct electric current easily
because the valence electrons in a metal can move freely among the
atoms.
Alloys – is a mixture made of two or more elements, at least
one of which is a metal.




Generally are stronger and less reactive than the pure metals from
which they are made.
Example: stainless steel = iron +carbon + nickel + chromium
Gold jewelry = gold+ cooper + silver
Assess your understanding
All bondings
All Bondings
 http://www.bsc2.ehb-
schweiz2.ch/Chemie/Simulationen%20Chemie/Bindung/Bi
ndung%20Hundeanalogie.htm
 Compare and contrast the following concepts:
Ionic bond
1.
2.
3.
4.
Covalent bond
1.
2.
3.
4.
Metallic Bond
1.
2.
3.
4.
END