Chemical Bonds The Formation of Compounds From Atoms Chapter 11

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Chemical BondsThe Formation of

Compounds From Atoms Chapter 11

Hein and Arena

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11.2Lewis Structures of Atoms

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Metals form cations and nonmetals form anions to attain a stable valence electron structure.

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This stable structure often consists of two s and six p electrons.These rearrangements occur by losing, gaining, or sharing electrons.

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• Na with the electron structure 1s22s22p63s1

has 1 valence electron.

The Lewis structure of an atom is a representation that shows the valenceelectrons for that atom.

• Fluorine with the electron structure 1s22s22p5

has 7 valence electrons

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The Lewis structure of an atom uses dots to show the valence electrons of atoms.

The number of dots equals the number of s and p electrons in the atom’s outermost shell.

BPaired electrons

Unpaired electron

Symbol of the element

2s22p1

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The number of dots equals the number of s and p electrons in the atom’s outermost shell.

S2s22p4

The Lewis structure of an atom uses dots to show the valence electrons of atoms.

911.4

Lewis Structures of the first 20 elements.

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11.3

The Ionic Bond

Transfer of Electrons FromOne Atom to Another

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The chemistry of many elements, especially the representative ones, is to attain the same outer electron structure as one of the noble gases.

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With the exception of helium, this structure consists of eight electrons in the outermost energy level.

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After sodium loses its 3s electron, it has attained the same electronic structure as neon.

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After chlorine gains a 3p electron, it has attained the same electronic structure as argon.

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Formation of NaCl

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The 3s electron of sodium transfers to the 3p orbital of chlorine.

Lewis representation of sodium chloride formation.

A sodium ion (Na+) and a chloride ion (Cl-) are formed.

The force holding Na+ and Cl- together is an ionic bond.

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Formation of MgCl2

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Two 3s electrons of magnesium transfer to the half-filled 3p orbitals of two chlorine atoms.A magnesium ion (Mg2+) and two chloride ions (Cl-) are formed.The forces holding Mg2+ and two Cl- together are ionic bonds.

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NaCl is made up of cubic crystals.In the crystal each sodium ion is surrounded by six chloride ions.

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In the crystal each chloride ion is surrounded by six sodium ions.

11.5

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The ratio of Na+ to Cl- is 1:1

There is no molecule of NaCl

11.5

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11.6

Electronegativity

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electronegativity: The relative attraction that an atom has for a pair of shared electrons in a covalent bond.

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• If the two atoms that constitute a covalent bond are identical, then there is equal sharing of electrons.

• This is called nonpolar covalent bonding.

• Ionic bonding and nonpolar covalent bonding represent two extremes.

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• If the two atoms that constitute a covalent bond are not identical, then there is unequal sharing of electrons.

• This is called polar covalent bonding.• One atom assumes a partial positive

charge and the other atom assumes a partial negative charge.–This charge difference is a result of the

unequal attractions the atoms have for their shared electron pair.

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:H Cl+ -

Shared electron pair.

:The shared electron pair is closer to chlorine than to hydrogen.

Partial positive charge on hydrogen.

Partial negative charge on chlorine.

Chlorine has a greater attraction for the shared electron pair than hydrogen.

Polar Covalent Bonding in HCl

The attractive force that an atom of an element has for shared electrons in a molecule or a polyatomic ion is known as its electronegativity.

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A scale of relative electronegativities was developed by Linus Pauling.

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Electronegativity decreases down a group for representative elements.Electronegativity generally increases left to right across a period.

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The electronegativities of the metals are low.The electronegativities of the nonmetals are high.

11.1

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The polarity of a bond is determined by the difference in electronegativity values of the atoms forming the bond.

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• If the electronegativity difference between two bonded atoms is greater than 1.7-1.9, the bond will be more ionic than covalent.

• If the electronegativity difference is greater than 2, the bond is strongly ionic.

• If the electronegativity difference is less than 1.5, the bond is strongly covalent.

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H H

Hydrogen Molecule

If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally.

The molecule is nonpolar covalent.

Electronegativity2.1


11.10

Electronegativity Difference = 0.0

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If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally.

Cl Cl

Chlorine Molecule



The molecule is nonpolar covalent.


11.10

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If the electronegativities are not the same, the bond is polar covalent and the electrons are shared unequally.

H Cl

Hydrogen Chloride Molecule



The molecule is polar covalent.

+ -


11.10

35Sodium Chloride

Na+ Cl-

If the electronegativities are very different, the bond is ionic and the electrons are transferred to the more electronegative atom.



The bond is ionic.No molecule exists.


11.10

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A dipole is a molecule that is electrically asymmetrical, causing it to be oppositely charged at two points.

A dipole can be written as + -

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An arrow can be used to indicate a dipole.

The arrow points to the negative end of the dipole.

H Cl H Br HO

H

Molecules of HCl, HBr and H2O are polar .

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A molecule containing different kinds of atoms may or may not be polar depending on its shape.

The carbon dioxide molecule is nonpolar because its carbon-oxygen dipoles cancel each other by acting in opposite directions.

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Relating Bond Type to Electronegativity Difference.

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11.7Lewis Structuresof Compounds

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In writing Lewis structures, the most important consideration for forming a stable compound is that the atoms attain a noble gas configuration.

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• The most difficult part of writing Lewis structures is determining the arrangement of the atoms in a molecule or an ion.

• In simple molecules with more than two atoms, one atom will be the central atom surrounded by the other atoms.

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Cl2O has two possible arrangements.

Cl-Cl-OThe two chlorines can be bonded to each other.

Cl-O-ClThe two chlorines can be bonded to oxygen.

Usually the single atom will be the central atom.

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Procedures for WritingLewis Structures

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Atom Group Valence Electrons

Cl 7A 7

H 1A 1

C 4A 4

N 5A 5

S 6A 6

P 5A 5

I 7A 7

Valence Electrons of Group A Elements

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Step 1. Obtain the total number of valence electrons to be used in the structure by adding the number of valence electrons in all the atoms in the molecule or ion.

–If you are writing the structure of an ion, add one electron for each negative charge or subtract one electron for each positive charge on the ion.

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Step 1. The total number of valence electrons is eight, two from the two hydrogen atoms and six from the oxygen atom.

Write the Lewis structure for H2O.

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Step 2. Write the skeletal arrangement of the atoms and connect them with a single covalent bond (two dots or one dash).– Hydrogen, which contains only one bonding

electron, can form only one covalent bond. – Oxygen atoms normally have a maximum

of two covalent bonds (two single bonds, or one double bond).

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Step 2. The two hydrogen atoms are connected to the oxygen atom. Write the skeletal structure:


Place two dots between the hydrogen and oxygen atoms to form the covalent bonds.

H O HorH OH

: : : :

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Step 3. Subtract two electrons for each single bond you used in Step 2 from the total number of electrons calculated in Step 1.–This gives you the net number of electrons

available for completing the structure.

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Step 3. Subtract the four electrons used in Step 2 from eight to obtain four electrons yet to be used.


H O H: :

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Step 4. Distribute pairs of electrons (pairs of dots) around each atom (except hydrogen) to give each atom a noble gas configuration.

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Step 4. Distribute the four remaining electrons in pairs around the oxygen atom. Hydrogen atoms cannot accommodate any more electrons.


These arrangements are Lewis structures because each atom has a noble gas electron structure.

H O HorH OH

: : : ::: ::

The shape of the molecule is not shown by the Lewis structure.

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Step 1. The total number of valence electrons is 16, four from the C atom and six from each O atom.

Write a Lewis structure for CO2.

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Step 2. The two O atoms are bonded to a central C atom. Write the skeletal structure and place two electrons between the C and each oxygen.

O C O: :


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Write a Lewis structure for CO2.Step 3. Subtract the four electrons used in Step 2 from 16 (the total number of valence electrons) to obtain 12 electrons yet to be used.

O C O: :

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O C O: :

Step 4. Distribute the 12 electrons (6 pairs) around the carbon and oxygen atoms. Three possibilities exist.

Many of the atoms in these structures do not have eight electrons around them.


O C O: : O C O: :: :::

:: ::: :

::

4 electrons

6electrons

6electrons

6electrons

:::

:: :

6electrons

I II III

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O C O: :: :::

::

Step 5. Remove one pair of unbonded electrons from each O atom in structure I and place one pair between each O and the C atom forming two double bonds.

O C O::: : ::

O C O::: : : :::::

Each atom now has 8 electrons around it.

Carbon is sharing 4 electron pairs.

double bond double bond

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11.8Complex Lewis Structures

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There are some molecules and polyatomic ions for which no single Lewis structure consistent with all characteristics and bonding information can be written.

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Step 1. The total number of valence electrons is 24, 5 from the nitrogen atom and 6 from each O atom, and 1 from the –1 charge.

Write a Lewis structure for NO2.-3NO .

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Step 2. The three O atoms are bonded to a central N atom. Write the skeletal structure and place two electrons between each pair of atoms.


O N O: :O:

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Step 3. Subtract the 6 electrons used in Step 2 from 24, the total number of valence electrons, to obtain 18 electrons yet to be placed.

O N O: :O:


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O N OO

Step 4. Distribute the 18 electrons around the N and O atoms.


:: : :::

:: :: ::

electron deficient

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O

:: : :::

:: :: ::O N O

Step 4. Since the extra electron present results in nitrate having a –1 charge, the ion is enclosed in brackets with a – charge.


-

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Step 5. One of the oxygen atoms has only 6 electrons. It does not have a noble gas structure. Move the unbonded pair of electrons from the N atom and place it between the N and the electron-deficient O atom, making a double bond.

NO

::O:: :

O:::

:-

::

electron deficient

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A molecule or ion that shows multiple correct Lewis structures exhibits resonance.


Step 5. There are three possible Lewis structures.

NO

::O:: :

O:

:

::

- :NO

::O:

:O::

::

-

Each Lewis structure is called a resonance structure.

NO

::O:: :

O:::

:-

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11.9Compounds Containing

Polyatomic Ions

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A polyatomic ion is a stable group of atoms that has either a positive or negative charge and behaves as a single unit in many chemical reactions.

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Sodium nitrate, NaNO3, contains one sodium ion and one nitrate ion.

sodium ion Na+ -3NOnitrate ion

NO

::O:: :

O:::

:

-Na +

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• The nitrate ion is a polyatomic ion composed of one nitrogen atom and three oxygen atoms.

NO

::O:: :

O:::

:

-Na +

• It has a charge of –1• One nitrogen and three oxygen atoms

have a total of 23 valence electrons.

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• The –1 charge on nitrate adds an additional valence electron for a total of 24.

NO

::O:: :

O:::

:

-Na +

• The additional valence electron comes from a sodium atom which becomes a sodium ion.

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• Sodium nitrate has both ionic and covalent bonds.

NO

::O:: :

O:::

:

-Na +

• Ionic bonds exist between the sodium ions and the carbonate ions.

covalent bond

covalent bond

covalent bond

ionic bond

• Covalent bonds are present between the carbon and oxygen atoms within the carbonate ion.

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• When sodium nitrate is dissolved in water the ionic bond breaks.

NO

::O:: :

O:::

:

-Na +

• The sodium ions and nitrate ions separate from each other forming separate sodium and nitrate ions.

NO

::O:: :

O::

::

-Na +

• The nitrate ion, which is held together by covalent bonds, remains as a unit.

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11.10

Molecular Shape

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The 3-dimensional arrangement of the atoms within a molecule is a significant feature in understanding molecular interactions.

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11.11The Valence Shell

Electron Pair (VSEPR) Model

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The VSEPR model is based on the idea that electron pairs will repel each other electrically and will seek to minimize this repulsion.

To accomplish this minimization, the electron pairs will be arranged as far apart as possible around a central atom.

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BeCl2 is a molecule with only two pairs of electrons around beryllium, its central atom. Its electrons are arranged 180o apart for maximum separation.

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• BF3 is a molecule with three pairs of electrons around boron, its central atom.

• Its electrons are arranged 120o apart for maximum separation.

• This arrangement of atoms is called trigonal planar.

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• CH4 is a molecule with four pairs of electrons around carbon, its central atom.

• An obvious choice for its atomic arrangement is a 90o angle between its atoms with all of its atoms in a single plane.

• However, since the molecule is 3-dimensional, the molecular structure is tetrahedral with a bond angle of 109.5o.

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Ball and stick models of methane, CH4, and carbon tetrachloride, CCl4.

11.13

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• Ammonia, NH3, has four electron pairs around nitrogen.

The arrangement of the electron pairs is tetrahedral.

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The shape of the NH3 molecule is pyramidal.

One of its electron pairs is a nonbonded (lone) pair.

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• Water has four electron pairs around oxygen.

The arrangement of electron pairs around oxygen is tetrahedral.

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The H2O molecule is bent.

Two of its electron pairs are nonbonded (lone) pairs.

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Structure Determination Using VSEPR

1. Draw the Lewis structure for the molecule.

2. Count the electron pairs and arrange them to minimize repulsions.

3. Determine the positions of the atoms.4. Name the molecular structure from

the position of the atoms.

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Chemical Bonds The Formation of Compounds From Atoms Chapter 11