Chemistry Day 27 Monday, November 18th – Tuesday,

November 19th, 2019

Do-Now: “Covalent Bonding CN B” 1. Write down today’s FLT 2. How can you tell if atoms will form an

ionic or a covalent bond? 3. Draw the electron dot diagram for Sulfur. 4. Use a pro-talk sentence frame to explain

the concept of covalent bonding. 5. What do you need to do to prepare for

the unit test? Take out your planner and ToC

FLT •  I will be able to describe the arrangement

of electrons covalent bonds using lewis dot structures by completing Covalent Bonding CN B

Standard HS-PS1-1:Usetheperiodictableasamodeltopredicttherelativepropertiesofelementsbasedonthepatternsofelectronsintheoutermostenergylevelofatoms.

The Nature of Covalent Bonding

Recall

Recall • Compounds are held together by

chemical bonds •  Ionic Bond = electrostatic attraction

due to the transfer of vse-s between a metal and nonmetal

• Covalent Bond = formed from the sharing of vse-s between nonmetals

Molecular vs. Ionic Compounds • NaCl • KBr • CO2 • HBr • NH3 • AlCl3

Single, Double, Triple Bonds

Single, Double, Triple Bonds • Covalent Bonds involve the sharing

of valence electrons •  These electrons form strong single,

double, or triple covalent bonds between the atoms

Single Bonds •  Single Bonds = Formed when atoms

share one pair of ve-s (2 total) •  Drawn as a line between atoms

•  Ex:

Single, Double, Triple • However, sometimes atoms can

share more than one pair of electrons.

• When this happens, double or triple bonds form.

Single, Double, Triple •  Double Bonds = atoms share two

pairs of ve-s (4 total) •  Stronger and shorter than single bonds •  Noted as two lines

Single, Double, Triple •  Triple Bonds = atoms share three

pairs of ve-s (6 total) •  Strongest, shortest, and most rigid •  Noted as three lines

Lewis Structures

Remind me… • How many valence electrons in… 1.  Li 2.  S 3.  Al 4.  Br 5.  Sr

Lewis Structures • We can

represent the covalent bonding patterns in atoms using Lewis structures

Rules for Drawing Lewis Structures 1.  Determine the total number of

ve-s ex/ PBr3

Rules for Drawing Lewis Structures 2. Place least EN atom in center and connect it to outer atoms w/ single bonds. Carbon is always in the center if present. Hydrogen is never in the center.

Rules for Drawing Lewis Structures 3. Use remainder of ve-s to fill in octets around outer atoms* *Atoms in 1s or 2s, like H, can only have 2 e-s

Rules for Drawing Lewis Structures 4. Place extra ve-s around the central atom

Rules for Drawing Lewis Structures 5. If you run out of ve-s, use pairs to form double or triple bonds ex/ C2H4

Things to Watch For: • Never add extra ve-s to H

Things to Watch For: • Add (-) charges and subtract (+)

charges from the total ve-s • Ex/ NO3

-

Things to Watch For: •  Indicate charges outside of

brackets • Ex/ NO3

-

Things to Watch For: • Resonance = multiple possible

structures (draw all with ßà) • Ex/ NO3

-

Extra Notes •  There are always exceptions

Extra Notes • Boron is a moron

Try This: SiBr4

1.  Find the total number valence electrons

2.  Draw the Lewis structure 3.  Compare with your neighbor

Molecular Geometry

The VSEPR Model •  VSEPR = Valence shell electron-pair

repulsion •  This model is used to predict the

geometries of molecules formed from nonmetals – The molecular structure is the 3D arrangement

of molecules in an atom •  The structure around a given atom is

mostly determined by minimizing electron-pair repulsions – Bonding and nonbonding e- pairs around an

atom will be placed as far apart as possible

Video Notes

The VSEPR Model Types of Molecular Structures: •  Linear structure – Molecule has a 180o

bond angle •  Ex/ BeCl2

The VSEPR Model Types of Molecular Structures: •  Trigonal Planar – Molecule has a planar

(flat) and triangular structure with 120o bond angles

•  Ex/

The VSEPR Model Types of Molecular Structures: •  Tetrahedral– Molecule has a bond angles of

109.5o. Occurs whenever four pairs of e-s

are present around an atom. •  Ex/

The VSEPR Model Types of Molecular Structures: •  Trigonal Bipyramidal •  Ex/

The VSEPR Model Types of Molecular Structures: •  Octahedral: •  Ex/

To determine the shape: 1. Count number of electron regions

around your central atom –  Any type of bond counts as one

region, and a lone pair also counts as one region

2. Determine how many of those regions are lone pairs

–  (memorize chart)

The VSEPR Model We can use the number of electron pairs to determine the probable molecular shape. Know these five:

The VSEPR Model We can use the number of electron pairs to determine the probable molecular shape. Know these five:

The VSEPR Model •  If there are no lone pairs, use the number

of bonding regions to predict the shape –  2 = linear –  3 = trigonal planar –  4 = tetrahedral –  5 = trigonal bipyramidal –  6 = octahedral

Try This: •  Draw the lewis structures, determine the

molecular geometry, AND sketch the shape for the following:

a.  PCl5

The VSEPR Model •  The presence of lone pairs count as

electron regions; however, they will change the molecular geometry

•  Ex/ H2O •  How many electron regions are around O?

The VSEPR Model •  Lone pairs require more room than bonded

(shared) pairs, and tend to compress the angles between the bonding pairs

The VSEPR Model •  To sum up:

– Determine the TOTAL number of electron regions around the central atom

– Determine how many of those are lone pairs

– Generally, any resonance structure can be used to determine geometry

The VSEPR Model

The VSEPR Model

The VSEPR Model

The VSEPR Model

The VSEPR Model

The VSEPR Model

Try This: •  Draw the lewis structures, determine the

molecular geometry, AND sketch the shape for the following:

a.  SO2

The VSEPR Model •  Remember: this is just a model •  While it generally provides us with an

accurate prediction, there are exceptions

CW1.   LewisStructureWS2.   StudyCh.123.   Finished?WorkonToCorstudy

quietly

Chemistry Day 28 Wednesday, November 20th – Thursday,

November 21st, 2019

Do-Now: “Bonding Quiz Review Do-Now” 1. Write down today’s FLT

2.  Distinguish between an ionic bond and a metallic bond.

3.  Identify two properties of ionic compounds 4.  Can calcium and phosphorous form an ionic

bond? Why or why not? 5. Write down the formula for strontium nitride 6. What is meant by the term “delocalized

electrons”? Be specific. 7.  Have you learned your geometries?

Take out your planner and ToC

Chemistry Day 29 Friday, November 22nd – Monday,

December 2nd, 2019

Do-Now: “Unit Exam Day Do-Now” 1.  Write down today’s FLT

2.  The element in group 7A & period 4 is ___. 3.  The ____ ____ are in group 2A and have ____

valence electrons. 4.  The formula for calcium chloride is ________

and it has a(n) _______ bond. 5.  If you are drawing a Lewis structure, when do

you form double or triple bonds? 6.  Which elements are naturally found as diatomic

molecules? 7.  Take 5 minutes to review for your unit test.

Take out your planner and ToC

FLT •  I will be able to demonstrate my

understanding of atoms and bonding by completing Unit 2 Exam

Standard HS-PS1-1:Usetheperiodictableasamodeltopredicttherelativepropertiesofelementsbasedonthepatternsofelectronsintheoutermostenergylevelofatoms.