Presentation on theme: "Covalent Compounds. Why do atoms bond? When a + nucleus attracts electrons of another atom Or oppositely charged ions attract( ionic bonds-metals and."— Presentation transcript:
Why do atoms bond? When a + nucleus attracts electrons of another atom Or oppositely charged ions attract( ionic bonds-metals and nonmetals) We know the octet rule- atoms tend to gain or lose or share e - in order to acquire a full set of 8 valence e -
Properties 1) Covalent compounds generally have much lower melting and boiling points than ionic compounds. 2) Covalent compounds are soft and squishy (compared to ionic compounds, anyway On the other hand, covalent compounds have these molecules which can very easily move around each other, because there are no bonds between them. As a result, covalent compounds are frequently flexible rather than hard.
More Properties 3) Covalent compounds tend to be more flammable than ionic compounds. 4) Covalent compounds don't conduct electricity in water. 5) Covalent compounds aren't usually very soluble in water.
More Properties There's a saying that, "Like dissolves like". This means that compounds tend to dissolve in other compounds that have similar properties (particularly polarity). Since water is a polar solvent and most covalent compounds are fairly nonpolar, many covalent compounds don't dissolve in water. Of course, this is a generalization and not set in stone - there are many covalent compounds that dissolve quite well in water.
Covalent Bond Chemical bond that results from sharing valence electrons. (Generally occurs when elements are close to each other on the periodic table and between nonmetallic elements) Molecule- a another name for covalent compound, is formed when two or more atoms bond covalently.
more Diatomic molecules- are molecules that occur in nature not as single atoms because the molecules formed are more stable than the individual atoms. There are 7 of these: Hydrogen (H 2 ), Nitrogen (N 2 ), Oxygen (O 2 ), Fluorine (F 2 ), Chlorine (Cl 2 ), Bromine (Br 2 ), and Iodine (I 2 ).
Naming Covalent Compounds Rule 1. The element with the lower group number is written first in the name; the element with the higher group number is written second in the name. Exception: when the compound contains oxygen and a halogen, the name of the halogen is the first word in the name. Rule 2. If both elements are in the same group, the element with the higher period number is written first in the name.
More Naming Rules Rule 3. The second element in the name is named as if it were an anion, i.e., by adding the suffix -ide to the name of the element. Rule 4. Greek prefixes are used to indicate the number of atoms of each nonmetal element in the chemical formula for the compound. Exception: if the compound contains one atom of the element that is written first in the name, the prefix "mono-" is not used.
Covalent Bonds When showing the bonding between atoms of covalent compounds, either a pair of dots or a line is the Lewis structure of a molecule. Lewis structures- use electron-dot diagrams to show how electrons are arranged in molecules
more Group 7A atoms have 7 valence e - and need 1 e- to fulfill the octet rule and become stable. So they will form a single covalent bond Do example Group 6A need 2 e- and will form 2 covalent bonds Group 5A need 3e- and will form 3 covalent bonds Group 4A will form 4 covalent bonds. Let’s practice!!!
Molecular Structures Several models can be used to represent a molecular structure: Molecular formula, structural formula, Lewis structure and ball and stick. Predicting the location of certain atoms: Hydrogen is always terminal Halogens usually terminal Elements with more than one atom are usually terminal Central atom- smallest electronegativity Find total number of electrons for bonding9 total valence electrons) Determine pairs (divide valence electrons by 2) Draw Bonds from central atom to other atoms Subtract bonded e- for total valence e- Place remaining e- around to complete octets If there are not enough electrons to give the atoms 8, except for hydrogen, make double and triple bonds. C,N,O and S can form double and triple bonds with same element and others.
Single and Multiple Bonds The sigma bond(σ)- single covalent bond (overlapping of the valence atomic orbitals resulting in the e- being in a bonding orbital between the two atoms) Multiple Covalent Bonds- atoms can attain a noble gas configuration by sharing more than one pair of electrons between two atoms.
Atoms can have Single - 1 pair of e- Double – 2 pair of e- Triple – 3 pair of e- Try CO 2, O 2, N 2
Electronegativity and Lewis Dot Structures When these atoms share electrons they are often not equally shared. This is due to electronegativity. Hydrogen and halogens usually bond to only one other atom and are usually on the outside of the molecule. The atom with the smallest electronegativity is often the central atom of the compound. When a molecule contains more atoms of one element than the others, these atoms often surround a central atom.
More A pi bond π is formed when parallel orbitals (p orbitals )overlap to share electrons Single bond- sigma bond Double bond- sigma and 1 pi bond Triple – sigma and 2 pi bonds
Strength of Covalent Bonds Covalent bonds involve attractive and repulsive forces Nuclei(+) and electrons(-) attract each other but Nuclei(+) repel nuclei(+) and (e-) repel (e-)
More There is a balance of attractive and repulsive forces When this balance is disrupted the bond could break Factors controlling Bond Strength 1. distance between nuclei= bond length
More Bond length Determined by size of atoms and number of shared electrons Size- length decreases when number of bonds increases Example- a triple bond ( 3 shared pairs of e-) is shorter than a double bond (2 shared pairs of electrons) Strength- shorter the length –stronger the bond
Bond Energy Energy change occurs when bonds are formed Bond Dissociation Energy- the amount of energy required to break a specific covalent bond- this is always a + value because energy is added to break a bond The sum of BDE for all bonds in a compound= the amt. of chemical PE available in one molecule of a compound BDE= the strength of the bond **as two atoms are bonded closer together the BDE increases Formation- energy released-exothermic Broken- energy added-endothermic
Examples of BDE and Bond Length BDE Bond Length F 2 F---F 159 kJ/mole short O 2 O=O 498 kJ/mole shorter N 2 NtripleN 945 kJ/mole shortest
Exceptions Some molecules and ions do not obey the octet rule. Reasons 1. some group has an odd number of valence e- and can’t form an octet Ex. NO 2 2. some groups form with few than 8 e- present. This is very reactive and will share an entire pair of e- which is a coordinate covalent bond. exBH 3 3. central atom contains more than 8 e- (expanded octet)- usually occurs with elements in period 3 or higher. Ex. PCl 5 or SF 6
Resonance Structures A condition when more than one valid Lewis structure can be written for a molecule or ion. This occurs when a molecule or ion has a combination of single and double bonds. Let’s try to draw the resonance structures of NO 3 - Now you try for homework SO 3, SO 2, O 3, NO 2 -
Molecular Shapes Valence Shell Electron Pair Repulsion VSEPR model Many reactions depend on the ability of two compounds contacting each other. The shape of the molecules determines whether they can get close enough to react. The repulsion of electron pairs in a molecule result in atoms existing at fixed angles to each other- Bond angle
Electronegativity and Polarity We know electronegativity( the attraction an atom has for electrons) increases as we go up and to the right of the periodic table. So Fluorine is the most electronegative element. If we have any of the diatomic molecules, we have identical atoms bonded together, their electronegativities are the same so they are nonpolar molecules. There is equal sharing of electrons
More Chemical bonds between elements are never completely ionic or covalent. When the difference in electronegativity between atoms increases it is more ionic When the difference in electronegativity between atoms is less it is more covalent Unequal sharing of electrons results in a Polar Covalent Bond
More Electronegativity difference Type of Bond % covalent and 50% ionic > 1.70Ionic < 1.70Covalent
More Molecules are either polar or non polar With polarity comes a partial (+) and partial (-) Show how partial charges spread out on H 2 O CCl 4
3 resonance structures
Dichromate Ion Lewis Structure Cr 2 O valence electrons