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Key Ideas

• Compounds&can&be&differentiated&by&their&chemical&and&physical&properties.&(3.1dd)&• Two&major&categories&of&compounds&are&ionic&and&molecular&(covalent)&compounds.&

(5.2g)&• Chemical&bonds&are&formed&when&valence&electrons&are&transferred&from&one&atom&

to&another&(ionic),&shared&between&atoms&(covalent),&mobile&within&a&metal&(metallic).&(5.2a)&

• In&a&multiple&covalent&bond,&more&than&one&pair&of&electrons&are&shared&between&two&atoms.&(5.2e)&

• Molecular&polarity&can&be&determined&by&the&shape&of&the&molecule&and&the&distribution&of&charge.&Symmetrical&(nonpolar)&molecules&include&CO2,&CH4,&and&diatomic&elements.&Asymmetrical&(polar)&molecules&include&HCl,&NH3,&and&H2O.&(5.2l)&

• When&an&atom&gains&one&or&more&electrons,&it&becomes&a&negative&ion&and&its&radius&increases.&When&an&atom&loses&one&or&more&electrons,&it&becomes&a&positive&ion&and&its&radius&decreases.&(5.2c)&

• When&a&bond&is&broken,&energy&is&absorbed.&When&a&bond&is&formed,&energy&is&released.&(5.2i)&

• Atoms&attain&a&stable&valence&electron&configuration&by&bonding&with&other&atoms.&Noble&gases&have&stable&valence&configurations&and&tend&not&to&bond.&(5.2b)&

• Physical&properties&of&substances&can&be&explained&in&terms&of&chemical&bonds&and&intermolecular&forces.&These&properties&include&conductivity,&malleability,&solubility,&hardness,&melting&point,&and&boiling&point.&(5.2n)&

• ElectronSdot&diagrams&(Lewis&structures)&can&represent&the&valence&electron&arrangement&in&elements,&compounds,&and&ions.&(5.2d)&

• Electronegativity&indicates&how&strongly&an&atom&of&an&element&attracts&electrons&in&a&chemical&bond.&Electronegativity&values&are&assigned&according&to&arbitrary&scales.&(5.2j)&

• The&electronegativity&difference&between&two&bonded&atoms&is&used&to&assess&the&degree&of&polarity&in&a&bond.&(5.2k)&

• Metals&tend&to&react&with&nonmetals&to&form&ionic&compounds.&Nonmetals&tend&to&react&with&other&nonmetals&to&form&molecular&(covalent)&compounds.&Ionic&compounds&contain&polyatomic&ions&have&both&ionic&and&covalent&bonding.&(5.2h)&

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UNIT 5 BONDING WWW.MRPALERMO.COM

LESSON 5.1 TYPES OF BONDS WWW.MRPALERMO.COM

What is a Bond?

•  Intramolecular force (between atoms) that hold one atom to another in a compound

•  The energy stored in a bond is potential energy

Why do atoms Bond?

•  Atoms bond together to get 8 valence electrons to become STABLE (Stable Octet)

•  Exception: Hydrogen can only have 2 (stable duet)

Forming a Bond

•  Energy is RELEASED (EXOTHERMIC) spontaneously

•  Forms a STABLE compound

Breaking a Bond

•  Energy is ABSORBED (ENDOTHERMIC) •  Stability decreases

•  Ex. Ripping two atoms apart requires ENERGY

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Remember BARF

Breaking bonds Absorbs (requires) energy Release of energy when bonds are Formed

Three Types of Bonds

1.  Ionic 2.  Covalent 3.  Metallic

Ionic Bonds

•  Occur between: -  METALS and NON-METALS -  TRANSFER of ELECTRONS from metal to

nonmetal to form a bond.

Example: sodium reacts with chlorine

electrostatic attraction

Properties of Ionic Compounds

•  Hard •  Crystalline structure •  HIGH Melt/Boiling Pt •  SOLUBLE in water •  Conduct ELECTRICITY only in solution (aq)

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Electrolyte

Compound that separates into ions in solution and is able to conduct electricity

Covalent Bonds

•  Occur between: -  NONMETALS and NONMETALS -  SHARING of ELECTRONS to obtain a full

valence shell (stable) -  Form MOLECULAR substances

**Trick to remember: SHARING is CARING

Why do they share instead of transfer?

•  Each element is not strong enough to remove (“steal”) an electron from the other.

Properties of Covalent Compounds

•  Soft •  Low melting and boiling pts (due to

weak attraction between molecules) •  Do not conduct electricity due to lack

of charged particles (ions)

Metallic Bonds

•  Between METAL atoms of the SAME element

•  Ex. Au atoms in a gold ring

Properties of Metals

•  High MP and high BP because bonds are strong.

•  Always capable of conducting electricity because of mobile electrons (freely flowing delocalized electrons)

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Sea of “mobile” electrons Example:

What type of bond is created in following: a. KBr b. HI c. NO d.  LiCl

POLYATOMIC IONS

•  Contain both Ionic and Covalent Bonds •  Located on Table E •  The atoms in polyatomic ions are held

together by covalent bonds •  Form IONIC COMPOUNDS with other

substances due to presence of ions

Ex. NH4Cl

EXAMPLE:

MgSO4

LESSON 5.2 BOND POLARITY WWW.MRPALERMO.COM

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Bond Polarity

•  Earth has 2 poles (north & south) •  Magnets also have 2 poles

•  Bonds may also have 2 poles (dipole) depending upon their electronegativities

TYPES OF COVALENT BONDS

Non-polar covalent bond

•  Equal sharing of electrons •  Electronegativity difference (E.N.D.)

between atoms 0 - 0.4 •  Usually between identical atoms •  Ex. H2

E neg 2.2 E neg 2.2

Polar Covalent Bond

•  Unequal sharing of electrons •  E.N.D. between atoms 0.5 – 1.7 •  One atom is slightly negative and one

atom is slightly positive. •  This is known as a dipole. •  Ex: HF

! Nonpolar

! Polar

!  Ionic

Comparing Ionic and Covalent Bonds How to determine the type of bond

1.  Determine if it is ionic or covalent 2.  If its covalent (NM & NM)than look up

the E.N.D. 3.  If E.N.D. is:

-greater than 0.4 its polar covalent -if it isn’t than it’s non polar covalent (0- 0.4)

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Summary of Bond Types

0.0 0.4 1.7-2.0

Covalent Ionic

Nonpolar Polar

Example: Ionic, Polar or NP

•  Na-S •  C-Br •  C-C •  H-O •  F-F •  K-O

LESSON 5.3 LEWIS DIAGRAMS FOR IONIC COMPOUNDS

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Lewis Dot Diagram: Ionic Compounds

•  Draw ion dot diagrams next to each other making sure that: -  The ion charges cancel out (add up to

zero) -  The opposite charged ions are next to

each other, and the like charged ions are as far away from each other as they can be.

Recall

•  Dot diagrams for positive ions (metals)

•  Dot diagrams for negative ions (nonmetals)

x

Example:

Dot diagram of NaCl Dot diagram of CaCl2

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Example: Chemical Formulas

•  Coefficient: in front of Formula •  Subscript: small # after an atom •  Subscript after ( ) multiply everything inside

by that # Example: 2Ca(NO3)2

# of atoms of each substance in the formula above: 2Ca, 4N, 12O

Example:

How many atoms of each substance are in the formula 3Na2SO4 ?

LESSON 5.4 LEWIS DIAGRAMS FOR COVALENT COMPOUNDS

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Lewis Dot diagrams: Covalent Compounds

1.  Write the element symbols next to each other (if more than two symbols write the UNIQUE symbol in the center)

2.  Count up the total number of valence electrons for all the elements

3.  Put 8 electrons around the central atom (if only two atoms pick one to place them around)

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4.  Distribute the remaining valence electrons to the other atoms equally until you run out

5.  Check to see if each atom has a complete valence shell (8 electrons except Hydrogen which has 2)

Cl

Cl

**A shared pair of electrons counts for both atoms **Each atom of Fluorine now has 8 electrons.

Total # of valence electrons for 2 chlorine atoms 7 x 2 = 14

Example H2O:

Total # of valence electrons for H2O = 2(1) + 6 = 8

Example F2:

Total # of valence electrons for F2 = 7 + 7= 14

Example: Draw dot diagram for HBr

6.  If all atoms do not have a full valence shell then you must ADD MULTIPLE BONDS (sharing of 2 or more PAIRS of electrons)

Example: CO2

How many electrons can be shared?

•  Single bond = sharing a pair (2) electrons

•  Double bond = sharing 2 pair (4) electrons

•  Triple bond = sharing 3 pair (6) electrons

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Example: Draw the dot diagram for O2 Example: How many electrons are shared in the C-H bond?

•  2 electrons (1pair) shared between

carbon and hydrogen

Example: How many electrons are shared in the C=C bond?

4 electrons (2 pair) shared between carbon and carbon

LESSON 5.5 MOLECULAR POLARITY

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Types of electron pairs

•  Bonding pairs - form bonds •  Lone pairs - nonbonding e-

Molecular Polarity depends on:

1.  Bond Polarity 2.  Shape of molecule

(Symmetrical vs. nonsymmetrical)

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Shape caused by: VSEPR

•  The Valence Shell Electron Pair Repulsion

•  Valence electrons are arranged as far from one another as possible to minimize the repulsion between them

(Like charges repel each other)

Polar Molecules

•  A molecule is polar if it: -  Contains POLAR BONDS -  Is ASYMMETRICAL (not symmetrical)

****If there are LONE PAIRS on central atom then the molecule is automatically POLAR

Non Polar Molecules

•  A molecule is nonpolar if it: -  Contains only NONPOLAR BONDS -  Is SYMMETRICAL

SHAPES OF POLAR MOLECULES

1. LINEAR

•  Must be asymmetrical

δ+ δ-

2. BENT

•  2 pair of lone pair electrons on central atom

δ-

δ+ δ+

.. ..

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3. PYRAMIDAL

•  1 lone pair of electrons on central atom

δ+

δ+

δ+

δ- SHAPES OF NONPOLAR MOLECULES

1. LINEAR

•  Must be SYMMETRICAL

δ+ δ+ δ- δ+ δ-

2. TETRAHEDRAL

•  Central atom bonded to 4 identical atoms

Summary of Shapes and Polarity

1 or 2 bonds only linear Polar if asymmetrical Nonpolar if symmetrical

2 bonded pairs and 2 lone pairs

bent polar

3 bonded pairs and 1 lone pair

pyramidal polar

4 bonded pairs and 0 lone pairs

tetrahedral Non-polar

If the central atom has ... the shape is Bond type

Example: Determine the molecular polarity and shape of CCl4

•  Nonpolar because its symmetrical •  Tetrahedral

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Example: Determine the molecular polarity and shape of H2S

•  Polar due to 2 pair of lone pair e- •  Bent

LESSON 5.6 INTERMOLECULAR FORCES

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WHAT TYPE OF FORCES HOLD MOLECULES TOGETHER?

Intermolecular Forces (IMF’s)

•  Weak forces of attraction BETWEEN molecules (covalent compounds)

Types of Intermolecular Forces

1.  Dispersion Forces 2.  Dipole-Dipole 3.  Hydrogen Bonding

Dispersion Forces

•  Weakest IMF •  Occurs between nonpolar

molecules •  Explains how nonpolar molecules

can exist in solid & liquid phase •  MORE ELECTRONS = GREATER

FORCE

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Example: H2 Dipole-Dipole

•  Between polar covalent molecules •  Partial negative end of dipole attracted

to partial positive end of another dipole •  The more polar the bond the greater the

IMF between the molecules

Example: HCl Hydrogen Bonding

•  Special case of dipole interaction. •  Strongest IMF •  Occurs between hydrogen of 1

molecule and F, O or N in another •  Remember

“H bonding is FON”

Example: NH3 Practice: What type of IMF occurs

between molecules of Cl2:

•  Dispersion force because Cl2 is nonpolar

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Practice: If a substance is a gas at STP what type of IMF does it have?

•  Dispersion Forces because they are weak and therefore they are easily overcome.