CHAPTER 4: ATOMIC THEORY AND BONDING

4.1 Pages 168-180

ATOMIC THEORY

 What is an atom? Answer: Smallest particle of any element (is a pure substance).

 What is a compound? Answer: A pure substance composed of two or more atoms.

For example: H2O (water)

What Makes up an Atom? Subatomic particles!!

3 Types of Subatomic Particles: 1.  Protons: 1+ (positive) electric charge 2.  Electrons: 1- (negative) electric charge 3.  Neutrons: uncharged, no electric charge

Where are they found? 1.  Nucleus: protons and neutrons 2.  Shells/energy levels: electrons exist

around the outside of the nucleus.

SIZE AND MASS OF AN ATOM Mass  Protons and neutrons are larger than electrons and

therefore have a greater mass.

Volume  Most of the volume of an atom is found in the region

outside of the nucleus where the electrons are found.  The nucleus makes up the smallest portion of

volume of an atom.

How small is the nucleus???? It would take about 10,000 – 100,000 nuclei to stretch

across the diameter of one atom!

HOW SMALL DID YOU SAY??!!

ATOMS AND THEIR NUCLEI  The Nucleus: contains both protons and neutrons

 General Rule: Atoms with < 30 protons tend to have an equal number of neutrons.

Heavy Atoms  Always more neutrons than protons  Stable, extra neutrons allow repelling protons more

distance between protons.

Very Heavy Atoms  Unstable  Repulsion between protons is very high due to extra

neutrons.

CHARGE OF THE NUCLEUS

 The nucleus always has a POSITIVE (+) charge.

Nuclear Charge  The charge on the nucleus  Found by counting the number of protons in an atom

or by its atomic # of a neutral atom.

Ok so...what is the nuclear charge of Nitrogen????

Nitrogen has = 7 protons charge is +7

Atomic # of a = # of protons = # of electrons neutral atom

HOW TO DETERMINE THE NUMBER OF SUBATOMIC PARTICLES OF AN ATOM:

  # of Protons = atomic number   # of Electrons = same as the number of protons   # of Neutrons = mass # - atomic number

For example:

REVIEW: ATOMIC NOTATION

THE PERIODIC TABLE REVIEW

 Each element is listed according to its atomic number row by row from left to right.

 Each row is called a PERIOD (same # of shells)

 Columns are called a FAMILY or GROUP (same # of valence electrons)

 Metals are on the LEFT and in the MIDDLE.

 Non-metals are in the upper RIGHT corner

 Metalloids form a staircase towards the right side.

 Transition metals CENTER from group 3 to 12

PERIODIC TABLE

IONS

  ION: when an atom either GAINS or LOSES electrons. Can have a (-) or (+) charge.

 CATIONS: metals, that have lost electrons to become positively (+) charged ions.

 MULTIVALENT: metals, that form ions in more than one way based on the type of reaction they undergo (eg Fe…what other metals are multivalent?)

 ANIONS: non-metals usually always gain electrons and form negatively (-) charged ions.

BOHR DIAGRAMS

 A modal invented by Niels Bohr in the 1900’s to show the number and arrangement of electrons in each shell/energy level.

 Rule: Each shell can only hold a certain number of electrons.

  1st shell 2 electrons   2nd shell 8 electrons   3rd shell 8 electrons   4th shell 18 electrons

HOW TO DRAW BOHR DIAGRAMS  Step 1: Using your periodic table determine the

# of protons, neutrons and electrons. Eg. Mg Protons = 12 Electrons = 12 Neutrons = atomic mass – atomic #

24 – 12 = 12

 Step 2: Draw the nucleus as a circle:

 Step 3: Put the # of protons in the nucleus and the # of neutrons below:

 Step 4: Place the electrons in orbits around the nucleus by drawing circles around the nucleus.

12 12

12 12

TRENDS WITHIN A PERIOD

PATTERNS OF ELECTRON ARRANGEMENT

1. Arrangement In Periods  The period # of an element equals the number of

occupied shells.

Eg: Look at your periodic table: Examine period 1: How many shells are there? How about in period 2?...period 3? Notice a trend???? (see page 175)

Also… As you move from left to right across any given

period within the periodic table each element has one more electron.

PATTERNS OF ELECTRON ARRANGEMENT

2. Arrangement in Groups  Each group has the same number of electrons in the

outermost shell or valence shell.

 Valence Electrons: electrons found in the valence shell.

 Look at your periodic table (or…pg 175) Notice how many valence electrons are found within groups 1, 2, 3, 4, etc. notice a trend?

STABLE OCTET

Stable Octet: the complete arrangement of 8 electrons in the outermost shell (a.k.a. valence shell).

 Look at the periodic table on pg 175:

Which group has a stable octet? Why?

Octet Rule: elements want to have the same number of electrons as the nearest noble gas because noble gases are very stable. 

FORMING COMPOUNDS

 Chemical Bond: When 2 atoms move closely together their valence electrons interact.

  Stability of an atom, ion or compound depends on its energy:

- Lower energy levels are more stable. (eg. Noble gases; He, Ne, Ar etc.)

3 WAYS ATOMS FORM A COMPOUND

  3 Ways an atom may acquire a valence shell like its nearest noble gas:

  Ionic Bonds: 1. Metal Atoms: always lose electrons to other atoms

non-metals forming cations (ionic bond).

2. Non-Metal Atoms: always gain electrons from other metal atoms forming anions (ionic bond).

 Covalent Bonds: 3. Atoms may share electrons (covalent bond of 2 non-metals).

Compounds are either ionic or covalent.

IONIC BONDING (METAL AND NON-METAL) When one or more electrons transfer from one atom to another. Ionic bonds contain (+) and (-)ions:

 Eg: magnesium oxide (MgO)

*Ionic compounds are made up of a metal and non-metal

HOW TO: BOHR DIAGRAM OF AN IONIC BOND  Recall: MgO

Mg O

2 + 2 -

COVALENT BONDING (NON-METALS)  When one or more unpaired electrons from each atom

pair together.

1. Covalent Compound: Contains at least 2 or more different elements Eg: HCl, H20 or CO2

2. Covalent Molecule: made up of 2 or more atoms: Eg: *HCl, H2O or O2

*Note: All compounds are molecules BUT not all molecules are compounds!

COVALENT BONDING (NON-METALS)

 Bonding pair: valence electrons involved in the covalent bond:

 Lone pair: electron pairs in the valence shell that are not involved in bonding.

LEWIS DIAGRAMS

Gilbert Lewis (1875-1946): American chemist invented a method to show bonding between atoms.

 Lewis Diagram: a diagram that shows only an atom’s valence electrons and chemical symbol:

HOW TO DRAW LEWIS DIAGRAMS  Step #1: Write down the element symbol

Eg: Nitrogen N

 Step #2: Determine how many electrons are found in the valence shell.

  (N = 7 electrons total, thus 5 valence electrons)

 Step #3: Draw dots representing valence electrons around element symbol. Draw dots singly until the 5th electron is reached.

N

LEWIS DIAGRAMS

Text: page 178

LEWIS DIAGRAMS OF IONS Rules for Drawing (+) and (-) Ions:  Step #1: metal or non-metal? Are valence electrons

removed or added?  Eg: NaCl Formation:

 Step #2: Place a bracket around each element and indicate the ions charge on the top right.

LEWIS DIAGRAMS: IONIC COMPOUNDS

Recall: involve a transfer (loss/gain) of valence electrons and contain both (+) and (-) ions.

 Metal and Non-metal bonding.

Eg. #1 Eg. #2

LEWIS DIAGRAM: COVALENT COMPOUNDS

Recall: Involves the sharing of valence electrons.  When 2 non-metals bond.

Notice how there is NO change in charge.

LEWIS DIAGRAMS: COVALENT MOLECULES Draw a Lewis Diagram of NH3 (ammonia): Step 1: Determine the number of valence electrons of

both N and H: (N=5, H=1) Step 2: Draw the chemical symbol for each element:

H N H

H

Step 3: Draw valence electrons by placing the correct # of dots.

Step 4: Draw single bonds by drawing a line to illustrate pairing of electrons.

DIATOMIC MOLECULES

= A pair of atoms that are joined by covalent bonds.  They form b/c as a 2 atom molecule they are more

stable (full octet) then on their own.  For example F2, O2, Cl2, Br2, I2 are all diatomic

molecules: