Chemical Bonds PDF

# Atomic Structure & Chemical Bonding ## Chemical Bonding ### Chapter 2: Chemical Bonding #### Table of Contents: 1. Lewis Dot Symbols 2. Bond Properties 3. Electronegativity 4. The Ionic Bond 5. The Covalent Bond 6. Intermolecular Forces 7. Exceptions of the Octet Rule ## Why do atoms of differe...

# Atomic Structure & Chemical Bonding ## Chemical Bonding ### Chapter 2: Chemical Bonding #### Table of Contents: 1. Lewis Dot Symbols 2. Bond Properties 3. Electronegativity 4. The Ionic Bond 5. The Covalent Bond 6. Intermolecular Forces 7. Exceptions of the Octet Rule ## Why do atoms of different elements react? Atoms of different elements combine in order to achieve a more stable electron configuration. Maximum stability results when an atom is iso-electronic with a Nobel gas. ## What is a chemical bond? The chemical bond is the force that holds atoms together molecules and ions in chemical compounds. ## Chemical Bonding ### 1. Lewis Dot Symbols When atoms interact to form a chemical bond, only their outer regions are in contact. For this reason, when we study chemical bonding, we are concerned primarily with the valence electrons of the atoms. To keep track of valence electrons in a chemical reaction, and to make sure that the total number of electrons dose not change, chemists use a system of dots devised by the American chemists Gilbert Newton Lewis. Lewis also formulated the Octet Rule: An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons. #### Lewis Dot System Consists of the symbol of an element and one dot for each valence electron in an atom of the element. - Except for helium, the number of valence electrons in each atom has the same as the group number of the element. - Elements in the same group have similar outer electron configurations and hence a similar Lewis dot symbols - We can not write simple Lewis dot symbols for transition metals, lanthanides and actinides. Why? ## Chemical Bonding ### 2. Bond Properties #### Bond Properties A. Bond Type - There are two types of bonds: - Ionic Bond: The electrostatic force that holds ions together in an ionic compound, (Na-Cl, Li-F). - Covalent Bond: A bond in which the electrons are shared by two atoms, (F-F, H-H, CH4). - Note that some references mention a third type of bonds, the Metallic bond. - Metallic Bond: Chemists use the electron-sea model to describe metallic bonding. The model proposes that the valence electrons of metal atoms move freely among the ions, forming a “sea” of delocalized electrons that hold the metal ions rigidly in place. B. Bond Degree - There are three degrees of bonds: - Single Bond: Two atoms are held together by one electron pair (bond), (H-H, H-F). - Double Bond: Two atoms share two pairs of electrons (2 bonds), (O=C=O). - Triple Bond: Two atoms share three pairs of electrons (3 bonds), (N=N). - Note: Double and triple bonds are called Multiple Bonds where two or more pairs of electrons are shared between two atoms. C. Bond Length - Is the distance between the nuclei of two boded atoms. - Note: for a given pair of atoms (such as C and C), triple bonds are shorter than double bonds which, in turn, are shorter than single bonds. D. Bond Energy - Is the energy required to break the bond. ## Chemical Bonding ### 3. Electronegativity - Electronegativity is a property that helps us distinguish an ionic bond from a covalent. Electronegativity is defined as: the ability of an atom to attract toward itself the electrons in a chemical band. - Elements with high electronegativity have a greater tendency to attract electrons than do elements with low electronegativity. - Electronegativity is related to electron affinity and ionization energy. - Thus, an atom such as fluorine, which has a high electron affinity (tends to pick up electrons easily) and a high ionization energy (does not lose electrons easily), has a high electronegativity. - On the other hand, sodium has a low electron affinity, a low ionization energy and a low electronegativity. - Electronegativity is a relative concept, meaning that an element’s electronegativity can be measured only in relation to the electronegativity of other elements. - Linus Pauling devised a method for calculating relative electronegativities of most elements. - In general, electronegativity increases from left to right across a period in the periodic table, as the metallic character of the elements decreases. - Within each group, electronegativity decreases with increasing atomic number, and increasing metallic character. - Note that, the transition metals do not follow these trends. - The most electronegative elements (the halogens, oxygen, nitrogen, and sulfur) are found in the upper right-hand corner of the periodic table, and the least electronegative element are clustered near the lower left-hand corner. ## Chemical Bonding ### 4. Ionic Bond - An ionic bond is the electrostatic force that holds ions together in an ionic compound. - For example the reaction between lithium and fluorine: - $Li + F \rightarrow Li^{+}F^{-}$ (or LiF) - $1s^{2}$ $1s^{2}2s^{2}2p^{6}$ - $1s^{2}2s^{1}$ $1s^{2}2s^{2}2p^{5}$ - The ionic bond in LiF is the electrostatic attraction between the positively charged lithium ion and the negatively charged fluoride ion. The compound itself is electrically neutral. - Imagine that this reaction occurs in separate steps: - (1) First the ionization of Li: - $Li \rightarrow Li^{+} + e^{-}$ - [(2)] The acceptance of an electron by F: - $F+ e^{-} \rightarrow F^{-}$ - (3) The two separate ions joining to form a LiF unit: - $Li^{+} + :F^{-} \rightarrow Li^{+}:F^{-}$. - **Exercise:** - Calcium burns in oxygen to form calcium oxide : - $2Ca(s) + O_{2}(g) \rightarrow 2CaO(s)$ - Write the steps of this reaction. - $Ca + O \rightarrow Ca^{2+} :O:^{2-}$ - $[Ar]4s^{2}$ $1s^{2}2s^{2}2p^{4}$ - $[Ar]$ $[Ne]$ ## Chemical Bonding ### 5. Covalent Bond - The Covalent Bond: A bond in which two electrons are shared by two atoms - The Covalent Compound: Are compounds that contain only covalent bonds. - **Notes:** - In a covalent bond, each electron in a shared pair is attracted to the nuclei of both atoms. - In a covalent bond, there isn’t any electrons transition from one atom to another, the two atoms share the electrons. - By sharing electrons in a covalent bond, the individual atoms can complete their octets. - There is only one unpaired electron in each Cl atom. Six of the seven Cl valence electrons are paired and they do not participate in covalent bond formation. - From each Cl atom, only one valence electron participate in the formation of Cl2. - The other, non-bonding electrons, are called lone pairs [ Pairs of electrons that are not involved in the covalent bond formation ] - **Types of Covalent Bonds:** - **Pure Covalent Bond (non-polar bond):** - In a molecules like H₂, F2, O2 and N₂, in which the atoms are identical (have the same electronegativity), the two electrons of the bond will be equally shared by the two atoms [ the electrons spend the same amount of time in the vicinity of each atom ]. - **Polar Covalent Bond (polar bond):** - In a molecules like HF, CH4, H₂O and HCN, in which the atoms are deferent (have deferent electro-negativity), the two electrons of the bond will not be equally shared by the two atoms [ ( the electrons doesn’t spend the same amount of time in the vicinity of each atom ]. - **Conclusion:** - ΔΕΝ between 1.7 and 3.3: mostly ionic - ΔΕΝ between 0.4 and 1.7: polar covalent - ΔΕΝ between 0.0 and 0.4: mostly pure covalent (non-polar) ## Chemical Bonding ### 6. Intermolecular Forces - The forces holding molecules together are generally called intermolecular forces. The energy required to break molecules apart is much smaller than a typical bond-energy. - There are 3 kinds of intermolecular forces - London dispersion force. - Dipole-dipole force. - Hydrogen-bonding force. - **Why are intermolecular forces important? ** - Intermolecular forces play important roles in determining the properties of a substances - They determine the phase of a substance at room temperature. - Strong Intermolecular Forces = Solids - Weak Intermolecular Forces = Gases #### London Dispersion Forces - Dispersion forces exist between all small covalent molecules - Electrons within the molecule momentarily shift away from one side of the molecule, setting up a momentary temporary dipole. - The attraction between the and of the temporary dipoles is called the dispersion force. This is an attractive force! - Dispersion forces are stronger for larger molecules and for straight chain molecules. - The attraction between temporary dipoles occurs between atoms and molecules only in non-polar substances - Tend to be stronger the larger the atom or molecule size (the more electrons are in the atom or molecule). - Relatively weak forces. #### Dipole-Dipole Forces - Attraction between polar molecules - Occurs when the partially positive end of one molecule attracts the partially negative end of another molecule. - Generally stronger than London dispersion forces. #### Hydrogen Bonding - (Diagram of water, ammonia, and ammonia in water) ## Chemical Bonding ### 7. Octet rule exceptions - **Lewis Structures that are Exceptions to the Octet Rule** - Exceptions to the Octet rule are: - Expanded octet. - Incomplete octets. - Molecules with an odd number of electrons. #### Expanded octet - Molecules in which an atom has more than an octet (this can occur with period 3 and higher elements...NEVER with period 2 elements). - **Phosphorus pentachloride:** - For $PCl_{5}$, 10 electrons are around the P central atom. - **Sulfur hexafluoride:** - For $SF_{6}$, 12 electrons are around the S central atom. #### Incomplete octets - Molecules in which an atom has less than an octet (generally occurs with B or Be and always with H). - **In $BF_{3}$, boron has an incomplete octet.** #### Molecules with an odd number of electrons - Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons. - **NO** - $N - 5e^{-}$ - $O - 6e^{-}$ - $11e^{-}$ # End of Chapter 2