4. Bonding SL

Describe the ionic bond as the electrostatic attraction between oppositely charged ions Describe how ions can be formed as a result of electron transfer

Ionic Bond- the attractive force between oppositely charged ions that were formed when electrons are lost from one atom and transferred to another.

            Cations- Positively Charged. Low Electronegativity

            Anions- Negatively Charged. High Electronegativity

Deduce which ions will be formed when elements in Groups 1, 2, and 3 lose electrons

Group 1, 2, and 3 elements form 1+, 2+, and 3+ ions respectively

Deduce which ions will be formed when elements in Groups 5, 6, and 7 gain electrons

Group 5, 6, and 7 elements form 3-, 2-, and 1- ions respectively

State that transition elements can form more than one ion

Transition elements can form many ions with different charges

            Iron can be 2+ or 3+

            Remember your examples from the Periodicity Unit as wel

Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values

Electronegativity- Relative measure of an atom’s ability to attract electrons

Difference in electronegativity in ionic bonds: Large

Ionic Bonds are usually between metals and non-metals

Elements on the left- Are not electronegative, so they lose electrons

Elements on the right- Are very electronegative, so they gain electrons

Elements in the middle- Are transition metals and lose electrons

State the formula of common polyatomic ions formed by non-metals in Periods 2 and 3

                                                                                     

Ammonium	NH4 +
Cyanide	CN -
Hydroxide	OH -
Ethanoate(Acetate)	C2­H­3­O2 -
Peroxide	O2  2-
Nitride	N 3-
Permanganate	MnO4 -
Carbonate	CO3 -
Phosphate	PO
Sulfite	SO3 2-
Sulfate	SO4 2-
Nitrite	NO2 -
Nitrate	NO3 -
Hypochlorite	ClO -
Chlorite	ClO2-
Chlorate	ClO3
Perchlorate	ClO4
Chromate	CrO4 2-
Dichromate	Cr2O7 2-
Mercuric	Hg +

Naming Rules

Group 1-3 Metal + Non-Metal- Cation + Anion-ide

                                                 (Example: MnO₂ is Magnesium Oxide)

Transition Metal + Non-Metal- Cation (Roman Numeral) + Anion-ide

                                                  (Example: Fe₂O₃ is Iron (III) Oxide)

Non-Metal + Non-Metal- Prefix-First Element + Prefix-SecondElement-ide

                                                   (Example: N₂O₅ is Dinitrogen Pentoxide)

                                 

Trick Table to Remembering Charges

1+ (AM)	3- (PN)	2- (CD COPSS)	1-
Ammonium Mercuric	Phosphate Nitride	Chromate Dichromate Carbonate Oxide Peroxide Sulfate Sulfite	Everything Else…

                              

Describe the lattice structure of ionic compounds

Giant Ionic Structures: Ionic bonds form a crystal-lattice structure. Electrostatic attraction between oppositely charged ions causes the negative ions to surround the positive ones and vice versa. The ions also have far reaching attractions to oppositely charged ions NOT immediately surrounding each.

Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei

Describe how the covalent bond is formed as a result of electron sharing.

Covalent bonds form between similarly charged atoms with high electronegativites that share pair(s) of electrons with one another when electron orbitals overlap.

Single bonds- Share one pair of electrons

Double bonds- Share two pairs of electrons

Triple bonds- Share three pairs of electrons

Dative covalent bond- SPECIAL EXCEPTION in which one compound donates both electrons to a bond.

            Example NH₃ + H⁺ à NH₄⁺  (because ammonia has a lone pair)

Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom.

 

State and explain the relationship between number of bonds, bond length, and bond strength

Multiple bonds are stronger than single bonds because more energy is needed create/brake the additional bond(s). They are shorter because the multiple electrons between the nuclei create a stronger negative charge, pulling the positive nuclei closer

Predict whether a compound of two elements would be covalent from the position of the elements on the Periodic table or from their electronegativity values

Difference between electronegativity in covalent bonds- Slight. Because neither compound is strong enough to completely steal one electron (as is the case with ionic bonds), pairs of electrons are shared.

Covalent bonds usually between non-metals

Predict the relative polarity of bonds from electronegativity values

Polarity- is the formation of slight charges (called dipoles) in a covalent bond because of electronegativity differences.

The more electronegative atom will form a slightly negative dipole, because electrons will like to “hang out” to the more attractive (electronegative) atom

The less electronegative atom will form a slightly positive dipole, because electrons will spend more time around the electronegative atom

SL OBJECTIVE: Predict the shape and bond angles of species with four, three, and two negative charge centers on the central atom using the valence shell electron pair repulsion (VSEPR) theory.

HL OBJECTIVE: Predict the shape and bond angles for species with five and six negative charge centers using the VSEPR theory

 

• This table will show you all the important info. The last column gives you an idea of the 3-D shape using examples. Those aren’t mandatory to know
• If you come across a double bond in a molecule, pretend it’s a single bond when trying to find shape and bond angle.
• REMEMBER! There is a difference between electron distribution and molecular geometry. Don’t mess these up!!!

Bond angle decreases slightly with each lone pair because lone pairs are slightly stronger than bonding pairs, so with each lone pair, the bond angle will close slightly by about 2 or 3 degrees.

Predict whether or not a molecule is polar from its molecular shape and bond polarities.

Polar molecules are typically asymmetrical. One polar bond sticking out of the molecule creates a polar charge

Nonpolar molecules are typically symmetrical. Polar bonds are pulling with equal force in opposite directions, creating no net effect

Describe and compare the structure and bonding in the three allotropes of carbon

Allotropes- different forms of an element that exist in the same physical state