IONS UNIT 1B.10
Ionic Bonding https://www.youtube.com/watch?
v=g_hwgHGeRhA https://www.youtube.com/watch
?v=WWc3k2723IM https://www.youtube.com/watch?v
=hiyTfhjeF_U

Bohr’s Model
 Electrons move like planets around the sun.
 In circular orbits at different levels.
 Amounts of energy separate one level from another.

III. The Bohr Model of the Atom
A.
Electrons of hydrogen circle the nucleus in orbits
1. orbits have a fixed amount of energy in the ground state
2. orbits are a fixed distance from the nucleus
3.
orbits furthest from the nucleus have the greatest energy

Niels Bohr
(1885 – 1962) Bohr Model of the Atom

Nucleus
Electron
Orbit
Bohr’s Model
Energy Levels

Bohr’s Model
Fifth
First
Fourth
Third
Second
}
 Further away from the nucleus means more energy.
 There is no “in between” energy
 Energy Levels
Nucleus

IONS
 An atom usually has a neutral charge. That means it has the same number of protons as electrons
 Remember, a proton has a positive charge and an electron has a negative charge
 ION – an atom that has lost or gained one or more electrons and has become charged either positively or negatively

Positive Ions
 When an atom LOSES electrons, it becomes more POSITIVE
– Why?
 If you are getting rid of negative particles
(electrons) but your number of positive particles (protons) are staying the same.
– In other words, you are subtracting negative numbers

Examples
 What would the charge be if:
 The neutral form of Gold (Au) lost 4 of its 79 electrons. It now has 79 protons and 75 electrons
 The neutral form of Mg lost 2 of its 12 electrons. It now has 12 protons and 10 electrons.

Negative Ions
 When an atom GAINS electrons it becomes more NEGATIVE
– Why?
 Electrons have a negative charge, so the more you have, the more negative you become

Representing Ions
 Ions are represented by placing a
“superscript” charge number next to the atomic symbol.
 Ex.
– O -2 = oxygen with a negative 2 charge
– K + = potassium with a positive 1 charge
– N -3 = nitrogen with a negative 3 charge
– And so on

A. Chemical Bond and valence electron
1.The electrons responsible for the chemical properties of atoms are those in the outer energy level: VALENCE ELECTRONS.
a.Valence electrons - The electrons in the outer energy level.
b.Inner electrons -those in the energy levels below.

Keeping Track of Electrons
2. Atoms in the same column a. Have the same outer electron configuration.
b. Have the same valence electrons.
c. Easily found by looking up the group number on the periodic table.
d. Group 2A - Be, Mg, Ca, etc.-
2 valence electrons

 List the number of valence shell electrons are in each of the elements in groups
 1-2 13-18

B. Electron Dot Diagrams: Lewis
Structures
1. A way of keeping track of valence electrons.
2. How to write them
3. Write the symbol.
4. Put one dot for each valence electron
5. Don’t pair up until they have to
X

The Electron Dot diagram for
Nitrogen
 Nitrogen has 5 valence electrons.
 First we write the symbol.
 Then add 1 electron at a time to each side.
 Until they are forced to pair up.
N

Write the electron dot diagram for
Na
Mg
C
O
F
Ne
He

C. Ion Formation in Representative
Elements
Group Gain or Lose Charge of Ion
1 (1A) lose 1 +1
2 (2A) lose 2 +2
13 (3A) lose 3 +3
14 (4A) lose or gain 4* +4,-4*
15 (5A) gain 3 -3
16 (6A) gain 2 -2
17 (7A) gain 1 -1

Name of ion
Cation keep the same name
Ca2+ Calcium
Al3+ Aluminum
Na+ Sodium
Anion ---ide
FFloride
P3phosphoride
O2oxide

C. Electron Configurations for
Cations
1. Metals lose electrons to attain noble gas configuration.
2. They make positive ions.
3. If we look at electron configuration it makes sense.
 Na 1level- 2e
 2 level -8e
 3 level -1e : 1 valence electron

Electron Dots For Cations
 Metals will have few valence electrons
Ca

Electron Dots For Cations
 Metals will have few valence electrons
 These will come off
Ca

Electron Dots For Cations
 Metals will have few valence electrons
 These will come off
 Forming positive ions
Ca +2

D. Electron Configurations for
Anions
1. Nonmetals gain electrons to attain noble gas configuration.
2. They make negative ions.
3. If we look at electron configuration it makes sense.
 S 1s 2 2s 2 2p 6 3s 2 3p 4 : 6 valence electrons
 S -2 1s 2 2s 2 2p 6 3s 2 3p 6 : noble gas configuration.

Electron Dots For Anions
 Nonmetals will have many valence
.electrons.
 They will gain electrons to fill outer shell.
P P -3

E. Stable Electron Configuration
1.All atoms react to achieve noble gas configuration.
2.Noble gases 8 electrons on last energy level.
3. 8 valence electrons .
4. Also called the octet rule.
Ar

 Na
 Mg
 P
 O
 F
 Cl
 K
Write the electron configuration diagram label as anion or cation

I. Chemical Bonds
 A. The force that holds two atoms together.
1. Why do atoms form bonds?
a. to acquire 8 electrons in the valence shell (like noble gases
2. How do atoms form bonds?
a. atoms may lose, gain or share electrons to get 8 in the valence shell

I. Properties of Ionic Compounds
Ionic_Bonds.asf
View Ionic video a. Crystalline structure.
b. A regular repeating arrangement of ions in the solid.
c. Structure is rigid.

II. Ionic Bonding
A. Anions and cations are held together by opposite charges.
B. Ionic compounds are called salts.
C. Simplest ratio is called the formula unit.
D.The bond is formed through the transfer of electrons.
E. Electrons are transferred to achieve noble gas configuration.

II.
Formation and Nature of Ionic Bonds
F . 1. atom “M” loses electron(s)  cation
2. atom “N” gains electron(s)  anion
3. cation and anion attract each other a. electrostatic attraction
4. the electrostatic force that holds the oppositely charged ions together is the ionic bond

Sodium loses an electron forming a (+) ion.
Chlorine gains an electron forming a ( - ) ion.
Electrostatic attraction between the (-) and (+) ion forms the ionic bond between sodium and chlorine

B. Ionic Compounds
1. compounds containing ionic bonds
2. types of ionic compounds a. oxides – metal + oxygen
Na
2
O, CaO, Al
2
O
3
, Fe
2
O
3 b. salt – metal + nonmetal
NaCl CaF AgCl KI c. binary compounds – two elements
1)all of the compounds in a and b are binary compounds

 D. monoatomic ion - one atom
1) ex. K + Fe 3+ O 2N 3-
E . polyatomic ion – ion with more than one atom that acts as a single ion
NO
3
OH SO
4
2NH
4
+

III. Names and Formulas
A.
Formulas for Ionic Compounds
1. vocabulary a. formula unit – simplest ratio of ions in a compound
1) ex. NaCl MgBr
2
AlCl b.
monoatomic ion - one atom
3
1) ex. K + Fe 3+ O 2N 3-

B . Formulas For Ionic Compounds
1. write formula for the cation first, then the anion
2. use subscripts to indicate number of ions (criss-cross the charges) a. sum of charges should equal 0 b. never change subscripts in polyatomic ions

Sodium Chloride Crystal

Ionic Bonding
Na Cl

Ionic Bonding: Lewis Structure
Na + Cl
-

Ionic Bonding
 All the electrons must be accounted for!
Ca P

Ca
Ionic Bonding
P

Ionic Bonding
Ca +2 P

Ionic Bonding
Ca +2
Ca
P

Ca +2
Ca
Ionic Bonding
P
-3

Ca +2
Ca
Ionic Bonding
P
-3
P

Ca +2
Ca +2
Ionic Bonding
P
-3
P

Ionic Bonding
Ca
Ca +2
Ca +2
P
-3
P

Ionic Bonding
Ca
Ca +2
Ca +2
P
-3
P

Ionic Bonding
Ca +2
Ca +2
Ca +2
P
-3
P
-3

Ionic Bonding
Ionic_Bonds.asf
Ca
3
P
2
Formula Unit


Shortcut
Ca +2 P -3
 Ca
3
P
2
 Crisscross the charges to become the subscript!!!!!!!!!!!!!!!!!!!!!!!!!

I. Properties of Ionic Compounds a. Crystalline structure.
b. A regular repeating arrangement of ions in the solid.
c. Structure is rigid.

Crystalline structure

Ionic Compounds
NaCl:
Ionic compounds consist of a lattice of positive and negative ions.
Lattice: three dimensional array of ions

Ionic properties d. Ions are strongly bonded- because of strong forces between ions they have
1. High melting points
2. high boiling point
3. high hardness scale
4. very rigid
5. very brittle

Ionic solids are brittle
+ + -
+ +
+ + -
+ +

Ionic solids are brittle
 Strong Repulsion breaks crystal apart.
+ +
+ + -
+ +

E. Conductivity
1.Conducting electricity is allowing charges to move.
2.In a solid, the ions are locked in place.
3. Ionic solids are insulators.
4. When melted, the ions can move around.
5. Melted ionic compounds conduct.
6. First get them to 800ºC.
7. ELECTROLYTE-Dissolved in water they conduct. (aqueous solution)

Building Ionic Compounds
Binary compound - metallic cation bonded to nonmetallic anion
1. Oxide- When a metal is ionically bonded to Oxygen
Salt - Metal + Non-metal

Polyatomic Ions
 An ion made up of two or more atoms bonded together that acts as a single unit with a net charge.
Build with tabs NH
4 ,
OH, Write the polyatomic ion and the charge it has.
Na(OH)- a balance compound
Mg +2 (OH ) –1
Mg(0H)
2 a balanced compound

ACTIVITY:
 MAKE TAB CUTOUTS OF THE VALENCE SHELL OF
 AL, O, Na, Cl, Mg, Ca
 Bond Al O


Mg O
Na Cl
 Build just one formula unitremember it doesn’t exist as one in nature – it is a crystal of many formula units where each – is surrounded by a +

Activity
 Demo- Burn to Create Mg+O=MgO
 Check the conductivity of:
 1. NaCl- solid
 2. NaCl-aqueous solution
 3. Distilled water
 4. Tap
 5. MgO-solid
 6. MgO-aqueous solution
 7. use magnifying glass to view NaCl

I. Metallic Bonds
 How atoms are held together in the solid form.
 Metals hold onto their valence electrons very weakly.
 Think of them as positive ions floating in a sea of electrons.

 http://www.youtube.com/watch?v=c4udBSZf
LHY

Sea of Electrons
 Electrons are free to move through the solid.
 Metals conduct electricity.
+ + + +
+ + + +
+ + + +

Metals are Malleable
 Hammered into shape (bend).
 Ductile - drawn into wires.

Malleable
+ + + +
+ + + +
+ + + +

Malleable
 Electrons allow atoms to slide by.
+ + + +
+ + + +
+ + + +

Video
 http://www.youtube.com/watch?v=Bjf9gMDP
47s

ALLOYS
 Alloys are solid solutions made by dissolving metals in other metals. They are prepared by melting the metals together and cooling the mixture.
 The properties of alloys differ from those of their component metals. For example stainless steel, an alloy of iron, carbon, chromium and nickel is stronger than iron and more resistant to corrosion.

Name of Alloy Composition
________________________________________
Sterling silver silver, copper
Brass copper, zinc, tin
Cast iron iron, carbon
Steel iron, carbon
Stainless steel iron, chromium, carbon, nickle
18 Carat gold gold, silver, copper
Pewter tin, copper, bismuth, antimony
Plumber’s solder lead, tin

ALLOY
 A combination of 2 metals

IV.
Metallic Bonds – Properties of Metals
A.
Metallic Bonds
1. valence electrons are delocalized a. free to move from atom to atom
2. bond is formed by the attraction of metal cations for the moving electrons
3. “ electron sea model” –atoms of metals contribute a “sea” of free moving electrons that move from one atom to another

Electron Sea Model for Metallic Bonds

Positive Ions Surrounded by Delocalized Electrons

B.
Properties of Metals
1. moderately high melting points
2. high boiling points
3. malleable a. can be hammered into sheets
4. ductile a. can be drawn into wire
5. good conductors of heat and electricity
6. luster (good reflectors of light)
7. hardness and strength varies a. greater in transition elements

C. Alloys
1. mixture of two or more elements with metallic properties
2. types a. substitutional – atoms of similar size (sterling silver, brass, pewter) b. interstitial – small holes in the crystal filled with smaller atoms
(carbon steel)