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Covalent Bonds (2 nonmetals) …atoms share e– to get a full valence shell C1s 2 2s 2 2p 2 F1s 2 2s 2 2p 5 *Both need 8 v.e – for a full outer shell (octet.

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Presentation on theme: "Covalent Bonds (2 nonmetals) …atoms share e– to get a full valence shell C1s 2 2s 2 2p 2 F1s 2 2s 2 2p 5 *Both need 8 v.e – for a full outer shell (octet."— Presentation transcript:

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3 Covalent Bonds (2 nonmetals) …atoms share e– to get a full valence shell C1s 2 2s 2 2p 2 F1s 2 2s 2 2p 5 *Both need 8 v.e – for a full outer shell (octet rule)!* o 4 valence e- 7 valence e- o x o o C x x x x x x F Bonding Review

4 Draw the Lewis dot structure for the following elements (write e- config first): Si O P B Ar Br 1s 2 2s 2 2p 6 3s 2 3p 3 1s 2 2s 2 2p 4 1s 2 2s 2 2p 1 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 1s 2 2s 2 2p 6 3s 2 3p 2 4 valence e- 6 valence e- 5 valence e- 3 valence e- 8 valence e- 7 valence e-

5 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Se Br Kr Rb Sr Te I Xe Cs Ba Notice any trends…? TRANSITION METALS The group # corresponds to the # of valence e –

6 C F F F F Let’s bond two F atoms together… Each F has 7 v. e – and each needs 1 more e – F F F F Now let’s bond C and F atoms together… C F F F F carbon tetrafluoride (CF 4 ) F2F2

7 Lewis Structures: 2D Structures NH 3 CH 2 O CO 2 SO 2 CH 4

8 1.Sum the # of valence electrons from all atoms Anions: add e – (CO 3 2- : add 2 e – ) Cation: subtract e – (NH 4 + : minus 1 e – ) 2.Predict the arrangement of the atoms Usually the first element is in the center (often C, never H) 3.Make a single bond (2 e – ) between each pair of atoms 4.Arrange remaining e – to satisfy octets (8 e – around each) Place electrons in pairs (lone pairs) Too few? Form multiple bonds between atoms: double bond (4 e – ) and triple bond (6 e – ) 5.Check your structure! All electrons have been used All atoms have 8e- Exceptions: Remember that H only needs 2e – ! Drawing Lewis Structures

9 Lewis Structure Practice CH 4 H 2 O NF 3 HBr OF 2 HCN NO 3 - CO 3 2- Draw a Lewis Structure for the following compounds: NHC

10 NHC

11 Lewis Structure Trends Here are some useful trends… C group Forms a combo of 4 bonds and no LP (Lone Pairs) i.e. CO 2 N group Forms a combo of 3 bonds and 1 LP i.e. NH 3 O group Forms a combo of 2 bonds and 2 LP i.e. CH 2 O F group (halogens) Forms 1 bond and 3 LP i.e. OF 2 Note that these are NOT always true!

12 Carbonite Carbonate? CO 3 2- CO 2 2-

13 Resonance Structures Resonance structures differ only in the position of the electrons The actual structure is a hybrid (average) of the resonance structures Technically NOT two single bonds and one double bond All 3 Oxygen atoms share the double bond 3 equal bonds (somewhere between a double and single) Arrow formalism: curved arrows show electron movement Show resonanceShow movement of e -

14 Predicting Molecular Shape: VSEPR (Valence Shell Electron Pair Repulsion) Electrons repel each other The molecule adopts a 3-D shape to keep the electrons (lone pairs and bonded e - ) as far apart as possible Different arrangements of bonds/lone pairs result in different shapes Shapes depend on # of bonds/lone pairs (“things”) and LP around the central atom

15 Selected Shapes and Geometries using VSEPR “Things”

16 Carbon Dioxide: CO 2 Two “things” (bonds or lone pairs) Linear geometry 0 LP → Linear Shape 180 o Bond angle COO Lewis Structure

17 C H H O Formaldehyde: CH 2 O Three “things” Trigonal planar geometry 0 LP → Trigonal planar shape 120° bond angles Lewis Structure

18 Sulfur Dioxide: SO 2 Three “things” Trigonal planar geometry 1 LP → Bent shape 120° bond angles Lewis Structure S O O B A A A

19 Methane: CH 4 Lewis Structure Four “things” (bonds/LP) Tetrahedral geometry 0 LP → Tetrahedral shape o bond angles

20 Ammonia: NH 3 Lewis Structure Four “things” (bonds/LP) Tetrahedral geometry 1 LP → Trigonal pyramid shape 107 o bond angles

21 Water: H 2 O Lewis Structure 4 “things” (bonds/LP) Tetrahedral Geometry 2 LP → Bent Shape o bond angle

22 Hydrogen Chloride: HCl Four “things” (bonds/LP) Tetrahedral geometry 3 LP → Linear Shape No Bond angle Lewis Structure ClH

23 A special note… For any molecule having only two atoms… e.g. N 2, CO, O 2, Cl 2, HBr, etc. Geometry = Linear Shape = Linear Bond Angle(s)? = None It is much like geometry… what is formed by connecting two points? …a line. NN HBr Cl OO

24 You will need to commit these to memory! “Things”

25 VSEPR Practice (w/o aid of yellow sheet) CO 2 G: S: Angle: ClO 2 - G: S: Angle: NO 2 - G: S: Angle: CH 3 COO - G: S: Angle: PBr 3 G: S: Angle: AsO 4 3- G: S: Angle:

26 Electronegativity and Bond Type The electronegativity difference between two elements helps predict what kind of bond they will form. Bond type Covalent  Polar covalent  Ionic Definition e- are evenly shared e- are unevenly shared e- are exchanged (gained or lost) Electronegativity difference ≤ 0.4  0.5 – 1.8  > 1.8

27 Electronegativity difference ≤ 0.4  0.5 – 1.8  > 1.8 Practice with Bond Types Bond type Covalent  Polar covalent  Ionic H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 Br 2.8 I 2.5 K 0.8 Ca 1.0 Sample Bonds NaCl Cl-Cl C-O C-H Electronegativity Difference 3.0 – 0.9 = – 3.0 = – 2.5 = – 2.1 = 0.4 Bond Type? Ionic Covalent Polar covalent Covalent

28 Dipole Moments and Polarity Arrow points toward partially “-” end Occurs in polar covalent bonds Uneven distribution of e - Atoms become partially charged Partially “+” charged end δ-δ- δ+δ+

29 Polarity Examples HCN CO 2 CO 3 2- CH 2 O SO 2 CH 4 CH 3 F C 3 H 8 CO NH 3 1.Check molecule for dipole moments (polar bonds) 2.When determining overall polarity, an imbalanced structure will likely be polar (at least partially) 3.Even with polar bonds, a balanced structure is non-polar overall 4. Any structure with lone pairs on the central atom is automatically polar! Try these with your neighbors… Polar Non-polar Polar Non-polar Polar Non-polar Polar


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