TOPIC 4 CHEMICAL BONDING AND STRUCTURE 4.2 COVALENT BONDING

Covalent compounds form by the sharing of electrons. ESSENTIAL IDEA Covalent compounds form by the sharing of electrons. NATURE OF SCIENCE (2.5) Looking for trends and discrepancies – compounds containing non-metals have different properties than compounds that contain non-metals and metals. NATURE OF SCIENCE (2.2) Use theories to explain natural phenomena – Lewis introduced a class of compounds which share electrons. Pauling used the idea of electronegativity to explain unequal sharing of electrons. .

UNDERSTANDING/KEY IDEA 4.2.A A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei.

ELECTRON SHARING When atoms of 2 non-metals react together, each is seeking to gain electrons in order to achieve the stable electron configuration of a noble gas. This tendency to form a stable arrangement of 8 electrons in the outer shell is referred to as the octet rule. The shared pair of electrons is concentrated in the region between the 2 positively charged nuclei. The electrostatic attraction between the 2 nuclei and the electrons constitutes the covalent bond.

UNDERSTANDING/KEY IDEA 4.2.B Single, double and triple covalent bonds involve one, two and three shared pairs of electrons, respectively.

UNDERSTANDING/KEY IDEA 4.2.C Bond length decreases and bond strength increases as the number of shared electrons increases.

BOND LENGTH AND BOND STRENGTH Triple bonds are stronger than double bonds which are stronger than single bonds. The strength of the bond is a measure of how much energy is required to break the bond. Triple bonds are shorter than double bonds which are shorter than single bonds. The number of shared electrons is greater in multiple bonds causing the electrostatic attraction to be stronger; therefore, causing the bonds to be shorter in length. A single bond contains only one sigma bond. A double bond contains one sigma and one pi bond. A triple bond contains one sigma and two pi bonds.

UNDERSTANDING/KEY IDEA 4.2.D Bond polarity results from the difference in electronegativities of the bonded atoms.

APPLICATION/SKILLS Be able to deduce the polar nature of a covalent bond from electronegativity values.

COVALENT CHARACTER You can predict covalent character by two ways: Position on the Periodic Table Electronegativity differences Covalent compounds tend to form between 2 non-metals. The closer together two elements are, the more covalent. Electronegativity Electronegativity values are given in Table 8 of the IB Data Booklet. Differences less than 1.8 are considered to be covalent.

POLARITY A bond that is unsymmetrical with respect to electron distribution is said to be polar. The term dipole means the bond has an area of positive charge and an area of negative charge. (delta symbol) The Pauling scale predicts polarity by using electronegativity differences. Values of 0.0-0.4 are considered non-polar covalent. Values between 0.4 and 1.8 are considered polar covalent. Values greater than and equal to 1.8 are considered ionic. Basically, the greater the electronegativity difference, the more polar the bond.

GUIDANCE Bond polarity can be shown either with partial charges, dipoles or vectors.

GUIDANCE Electronegativity values are given in the data booklet on page 8.

Aims Aim 3: Use naming conventions to name covalently bonded compounds.

Diatomic Elements Gases that exist as diatomic molecules are H2, F2, N2, O2, Cl2, Br2, I2 They are simply given their elements name. Exist this way only when not in compounds

Learning Check Use the name of the element to name the following diatomic molecules. H2 hydrogen N2 nitrogen Cl2 _______________ O2 _______________ I2 _______________

Naming Covalent Compounds Two nonmetals Name each element End the last element in -ide Add prefixes to show more than 1 atom Prefixes mon 1 hexa 6 di 2 hepta 7 tri 3 octa 8 tetra 4 nona 9 penta 5 deca 10

Learning check Fill in the blanks to complete the following names of covalent compounds. CO carbon ______oxide CO2 carbon _______________ PCl3 phosphorus _______chloride CCl4 carbon ________chloride N2O _____nitrogen _____oxide

Dinitrogen Hexafluoride Formula Names N2F6 Dinitrogen Hexafluoride CO2 Carbon Dioxide SiF4 Silicon Tetrafluoride CBr4 Carbon Tetrabromide NCl3 Nitrogen Trichloride P2S3 Diphosphorous Trisulfide CO Carbon Monoxide NO2 Nitrogen Dioxide SF2 Sulfur Difluoride PF5 Phosphorous Pentafluoride SO2 Sulfur Dioxide NO Nitrogen Monoxide CCl4 carbon tetrachloride P2O5 diphosphorus pentoxid

Learning Check A. P2O5 1) phosphorus oxide 2) phosphorus pentoxide 3) diphosphorus pentoxide B. Cl2O7 1) dichlorine heptoxide 2) dichlorine oxide 3) chlorine heptoxide C. Cl2 1) chlorine 2) dichlorine 3) dichloride

Citations International Baccalaureate Organization. Chemistry Guide, First assessment 2016. Updated 2015. Brown, Catrin, and Mike Ford. Higher Level Chemistry. 2nd ed. N.p.: Pearson Baccalaureate, 2014. Print. Most of the information found in this power point comes directly from this textbook. The power point has been made to directly complement the Higher Level Chemistry textbook by Catrin and Brown and is used for direct instructional purposes only.