Chapter 11 - Part 1 Bonding

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Foundations of College Chemistry, 14th Ed.

Morris Hein and Susan Arena

Atoms in Vitamin C (ascorbic acid) bond in a specific orientation which defines the shape of the molecule. The molecules pack

in a crystal, photographed above in a polarized micrograph (magnification 200x).

11 Chemical Bonds: The Formation of

Compounds from Atoms

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Metals, Nonmetals, and Metalloids

Periodic trends permit us to predict chemical properties and reactivity of the elements.

Metal properties: usually lustrous, malleable and good heat/electrical conductors.

Nonmetal properties: usually not lustrous, brittle and poor heat/electrical conductors.

Metalloid properties: can have some properties of either metals or nonmetals.

Tend to lose electrons to form cations.

Tend to gain electrons to form anions.

Periodic Trends in Atomic Properties

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Classification of Metals, Nonmetals, and Metalloids

Hydrogen is an exception – properties more like a nonmetal.

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Periodic Properties

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Atomic Radius

General trend: increases down a group and decreases left to right across a period.

Down a Group:

Additional n quantum levels are added; electrons are farther from the nucleus, so size increases.

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Ionization Energy

Energy required to remove an electron from an atom. Na + ionization energy Na+ + e–

1st Ionization Energy: energy needed to remove

the first electron.

2nd Ionization Energy: energy needed to remove

the second electron.

Successive ionizations always increase in energy as the remaining electrons feel stronger attraction

to the resulting cation produced.

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1st Ionization Energies for the Main Group Elements

General trend: generally decreases down a group and increases left to right across a period.

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Ionization Energies for Selected Elements

Extreme amounts of energy needed to remove electrons from noble gas configurations (in cyan).

Note: nonmetals generally have large ionization energies. Nonmetals usually gain electrons to form anions.

reserved.

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Valence Electrons

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Valence (outermost) electrons in an atom are responsible for the formation of chemical bonds between atoms.

Lewis structures are a simple way of representing atoms and their valence electrons.

Lewis structures use dots to represent the valence electrons of an atom.

Example: Boron

Electron Configuration: [He] 2s22p1

Lewis structure of an atom

Lewis Structures of Atoms

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Lewis Dot Structures of the First 20 Elements

Determining the Valence Electrons for an Atom:

For main group elements, the group number gives the number of valence electrons.

Example: Chlorine

reserved.

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Isoelectronic

• _{Isoelectronic species have the same electron}

configuration.

• _{Atoms tend to gain or lose electrons to}

become isoelectronic with noble gases

• _{Ne 1s}

2

_2s

2

_2p

6

• _Na

+1

_1s

2

_2s

2

_2p

6

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Lewis Electron Dot Structures

• _{For ions}

–

_{Add or subtract dots for electrons gained or}

lost to form ion.

–

_{For O}

2-

_–

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Bonding

• _{Atoms like to have a full outer shell and will}

gain lose or share electrons to achieve a full

outer shell

• _{Representative elements gain, lose, or share}

electrons to have 8 electrons in their outer

shell corresponding to full s and p orbitals.

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Two kinds of bonds

• _Ionic

–

_{atoms gain or lose electrons to form octet}

–

_{ions held together by electrostatic forces}

• _Covalent

–

_{atoms share electrons to form octet}

–

_{atoms held together by shared electron}

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Ion Formation

Cations

• _Na



Na

+

_{+ 1e}



• _Mg



Mg

+2

_{+ 2e}



Anions

• _{Cl + 1 e}





_Cl

1

• _{S + 2 e}





_S

2

These elements tend to lose electrons to gain an octet

Note: losing an electron always costs energy -- but sometimes this lost of energy can be compensated by the strong electrostatic energy gained when a cation and an anion combined.

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Ionic Bonds

• _{Bonds formed by the interaction of ions}

and the strong electrostatic forces that hold

them together.

• _{Ions group together in ratios which balance}

their positive and negative charges

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In the solid state, NaCl packs in an ordered

three dimensional cubic array. Each Na+ is surrounded

by 6 Cl– anions. In turn, each Cl– is surrounded

by 6 Na+ cations.

Lattice Diagram for NaCl

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Six electrons need to be transferred, three each from two different Al atoms, to balance transfer of

two electrons each to three O atoms.

More than one electron can be transferred between a metal and a nonmetal to form ionic compounds.

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In almost all stable compounds of main group elements, all atoms attempt to attain noble gas configuration. Metals lose electrons to do so; nonmetals gain electrons.

Predicting the Formula of Ionic Compounds:

Ionic compounds are charge neutral.

(The overall + and – charges must balance!)

Example: What is the formula containing Ba and S? Ba is in Group 2, would like to form Ba2+

S is in Group 6, would like to form S

2-Charge balance of Ba2+ and S2- = BaS

Predicting Formulas of Ionic Compounds

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Cations are smaller than the parent atom while anions are larger than the parent atom.

Relative radii of Na and Cl atoms and their corresponding ions.

In NaCl, Na+ is much smaller than Na, as there is greater

attraction between the nucleus and the 10 electrons remaining relative to the parent atom.

Cl– is larger than Cl because the attraction between the

nucleus is reduced by adding an extra electron to make an anion.

Cation/Anion Size

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Covalent Bonds

• _{Bonds where atoms share electrons to}

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Molecules are discrete units held together by covalent bonds.

Covalent bond: pair of electrons shared between 2 atoms

Formation of the H₂ molecule from two hydrogen atoms. The 1s orbitals on each atom overlap, forming a covalent bond between the atoms.

Each atom contributes one electron to the bond. Example: the covalent bond in H₂

Covalent Bonds

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The bonding in Cl₂

Overlap of a 3p orbital on each chlorine gives rise to a bond between chlorine atoms.

Each atom contributes one electron to the bond.

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Single bonds are typically formed between hydrogen or halogens because each element needs 1 electron

to achieve a noble gas configuration.

Typically, shared electrons between atoms are replaced with a line (--).

Atoms, to be more stable, can also form 2 or 3 covalent bonds as needed.

Covalent Bonds Between Other Elements

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Ionic and covalent bonding represent two extremes. Electrons are either fully transferred between

atoms (ionic) or shared (covalent).

Realize bonding in between the extremes also exists (i.e. polar covalent bonds).

Purely ionic character Purely covalent character Polar covalent

Electron sharing Electron transfer

Bonding: A Continuum

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Electronegativity: a measure of the ability of an atom to attract electrons in a covalent bond.

Example

Chlorine is more electronegative than hydrogen. The electrons shared between H and Cl remain closer

to Cl than H as a result. This gives the Cl a partial (δ) negative charge. H then has a partial positive charge.

Electronegativity

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Electronegativity

• _{A measure of the relative tendency of an}

atom to attract electrons to itself when it is

bonded to another atom.

• _{“Electron Greed”}

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A scale of relative electronegativities was developed by Pauling.

The most electronegative element, F, is assigned a value of 4.0.

The higher the electronegativity, the stronger an atom attracts electrons in a covalent bond.

Bond polarity is determined by the difference in

electronegativity of the two atoms sharing electrons.

Electronegativity Scale

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Bond polarity is determined by the difference in

electronegativity of the two atoms sharing electrons. If the atoms in the covalent bond are the same, the bond

is nonpolar and the electrons are shared equally. If the difference in electronegativities is > 2,

the bonding is considered ionic.

If the difference in electronegativties is < 2, the bonding is considered polar covalent (unequal sharing).

Bonding Continuum

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Ionic Bonds

An ionic bond

• _{occurs between metal and nonmetal ions} • _{is a result of electron transfer}

• _{has a large electronegativity difference (1.8 or more)}

Examples:

Atoms Electronegativity Type of Bond Difference

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Nonpolar covalent bonds have very small or no difference in electronegativities between the two atoms in the bond.

Electrons are shared equally between atoms.

Examples:

C-S bonds: electronegativity difference = 2.5 - 2.5 = 0

N-Cl bonds: electronegativity difference = 3.0 - 3.0 = 0

Bonding Continuum: Nonpolar Bonds

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Polar covalent bonds occur when two atoms share electrons unequally.

Examples:

P-O bonds: electronegativity difference = 3.5 -2.1 = 1.4

N-C bonds: electronegativity difference = 3.0 -2.5 = 0.5

P O

N C

Bonding Continuum:

Polar Covalent Bonds

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Polar Bonds

• _{Bonds in which electrons are not shared}

equally due to the electronegativity

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A molecule can contain polar bonds (if the two atoms are different) but still be nonpolar, depending on the

geometry of the molecule. Example:

Though the molecule contains polar C=O bonds, the molecule is linear, so the C-O dipoles cancel each other.

Symmetric arrangements of polar bonds result in nonpolar molecules.

Carbon dioxide

O = C = O

Asymmetric arrangements of polar bonds result in polar molecules.

N

H

H H

Polar Bonds and Polar Molecules

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Molecules which are electronically asymmetric have a permanent dipole, which means the molecule is polarized

with negatively and positively charged regions.

A dipole may also be written as a vector arrow, with the the arrow pointing toward the

negatively charged end of the molecule.

Dipole Moments

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Polar Molecules

• _{Dipole - A molecule such as HF which has a}

positive and a negative end. This dipolar

character is often represented by an arrow

pointing towards the negative charge.

• _{Dipole moments – the measure of the net}

molecular polarity

• _{Measured in units of Debyes (D) = Qr (charge x}

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51 Use electronegativity differences to classify each of

the following bonds as nonpolar covalent (NP), polar covalent (P), or ionic (I):

A bond between A. K and N

B. N and O C. Cl and Cl D. H and Cl

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52 Atoms in Electronegativity Type of

Bond Difference Bond

A. K and N 2.2 ionic (I)

B. N and O 0.5 polar covalent (P)

C. Cl and Cl 0.0 nonpolar covalent (NP) D. H and Cl 0.9 polar covalent (P)