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Home , Van der Waals force

Chapter 8 “Covalent Bonding” Bonds are… … that hold groups of together Section 8.1 – Molecular Compounds and make them function as a unit. Two types: 1) Ionic bonds – transfer of electrons (gained or lost; • OBJECTIVES: makes formula unit) – Distinguish between the melting points 2) Covalent bonds – sharing of electrons. The and boiling points of molecular resulting particle is called a “” compounds and ionic compounds. . Many elements found in nature are in the form of :

– Describe the information provided by a  a neutral group of atoms joined together by covalent bonds. molecular formula. . For example, air contains molecules, consisting of two oxygen atoms joined covalently

. Called a “diatomic molecule” (O2)

How does H2 form? Covalent bonding (diatomic hydrogen molecule) • Nonmetals hold on to their valence electrons. • The nuclei repel each other, since they both have a positive charge • They don’t give away electrons to bond. (like charges repel). – But still want noble configuration. • But, the nuclei are attracted to the • Get it by sharing valence electrons with each other = electrons covalent bonding • They share the electrons, and this is • By sharing, both atoms get to count the electrons toward called a “”, and involves only NONMETALS! a configuration.

 Fluorine has + + seven valence e- 8 Valence electrons  A second F F also has seven + +  By sharing e-…  …both end with full orbitals

Covalent bonding • Compounds that are bonded covalently (like in , or Molecular Compounds dioxide) are called molecular compounds • Thus, molecular compounds tend to be or liquids at • Molecular compounds tend to have relatively lower melting and room temperature boiling points than ionic compounds – this is not as strong a bond as ionic – Ionic compounds were . Fluorine has seven • A molecular compound has a molecular formula: valence e- – Shows how many atoms of each element a molecule . A second atom also has seven contains • 8 Valence . By sharing The formula for water is written as H2O electrons… – The subscript “2” behind hydrogen means there are 2 electrons F F . …both end with atoms of hydrogen; if there is only one atom, the full orbitals subscript 1 is omitted • Molecular formulas do not tell any information about the structure (the arrangement of the various atoms).

1 - Page 215 Section 8.2 The Nature of Covalent Bonding These are some of the 3. The ball and stick model is different ways to represent the BEST, because it shows • OBJECTIVES: ammonia: a 3-dimensional arrangement. – Describe how electrons are shared to form covalent bonds, and 1. The molecular identify exceptions to the octet rule. formula shows – Demonstrate how electron dot structures represent shared how many atoms electrons. of each element – Describe how atoms form double or triple covalent bonds. are present – Distinguish between a covalent bond and a , and describe how the strength of a covalent bond is 2. The structural related to its bond dissociation . formula ALSO – Describe how oxygen atoms are bonded in ozone. shows the arrangement of these atoms!

A Single Covalent Bond is... Water  Each hydrogen has 1 valence electron • A sharing of two valence electrons. H - Each hydrogen wants 1 more • Only nonmetals and hydrogen.  The oxygen has 6 valence electrons • Different from an ionic bond because they actually form molecules. - The oxygen wants 2 more • Two specific atoms are joined.  They share to make each other complete Note the two • In an ionic , you can’t tell which atom the electrons moved O “unshared pairs” from or to of electrons • Ionic compounds H O organize in a characteristic • It takes 2 hydrogens lattice of alternating to provide enough e- positive and negative • Every atom has full H There are 7 diatomic ions, repeated over energy levels elements, KNOW THEM: and over. H2 N2 O2 F2 Cl2 Br2 I2

Lewis Structures How to Draw a Lewis Structure • add up the valence e- from each element to get the total • a Lewis structure represents a chemical formula: the nuclei number of e- that can be used for bonding purposes and inner-shell e- are represented by the element’s atomic • when drawing a Lewis structure for a molecule with more than symbol, and covalent bonds are represented by pairs of dots two atoms, use the following guidelines for arrangement – hydrogen and halogen atoms usually bond to only one other atom in a or dashes molecule and are usually on the outside or end of a molecule – one pair is a single covalent bond (2 e-) that is shared – the atom with the smallest electronegativity is often the central atom – when a molecule contains more atoms of one element than the – two pair are a double covalent bond (4 e-) that is shared others, these atoms often surround the central atom – three pair are a triple covalent bond (6 e-) that is shared • using trial and error and only the number of valence e- - - allowed, place dots so that every element satisfies the • valence e are the e in the highest occupied energy level for octet rule (except hydrogen, which only needs 2 e-) any given element • dots that are shared should be changed to dashes, and • there can also be an unshared pair of e-; it is not involved in unshared pairs left as dots covalent bonding, but instead belongs exclusively to one • when more than one Lewis structure can be drawn for a given atom molecule, the molecule is said to be a resonance structure

2 Resonance Structures Bond Dissociation ... Neither ozone structure is correct, it is actually a hybrid of • The total energy required to break the two. To show it, draw all varieties possible, and join them the bond between 2 covalently with a double-headed arrow. Note the different location of the double bond bonded atoms • Occur when more than one valid Lewis structure can be written for • High dissociation energy usually a particular molecule (due to position of double bond) • Polyatomic ions – note the different positions of the double bond means the chemical is relatively 2- • Resonance in a carbonate ion (CO3 ): unreactive, because it takes a lot of energy to break it down.

Section 8.3 – Bonding Theories VSEPR stands for... • OBJECTIVES: • Valence Shell Electron Pair Repulsion – Describe the relationship between atomic and • Predicts the three dimensional shape of molecules. molecular orbitals. • The name tells you the theory: – Describe how VSEPR theory helps predict the shapes – Valence shell = outside electrons. of molecules. – Electron Pair repulsion = electron pairs try to get as far away as possible from each other. Molecular Orbitals are... • Can determine the angles of bonds. • The model for covalent bonding assumes the orbitals are • Based on the number of pairs of valence electrons, those of the individual atoms = both bonded and unbonded. • Orbitals that apply to the overall molecule, due to atomic • Unbonded pair also called lone pair. orbital overlap are the molecular orbitals – A bonding orbital is a that can be occupied by two electrons of a covalent bond

Tetrahedral and Other Shapes Section 8.4 Polar Bonds and Molecules • Methane (CH4) has a tetrahedral shape = 109.5° • OBJECTIVES: • Carbon dioxide (CO ) has a linear shape = 180o 2 – Describe how electronegativity (EN) values • Ammonia (NH3) has a determine the distribution of charge in a polar trigonal pyramidal H 109.5º molecule. shape = 107o – Describe what happens to polar molecules when they are placed between oppositely • Water (H2O) has a charged metal plates. bent shape = 105o C – Evaluate the strength of intermolecular H H attractions compared with the strength of ionic and covalent bonds. H – Identify the reason why network solids have high melting points.

3 Bond Polarity Bond Polarity • Covalent bonding means shared electrons • Refer to periodic table for EN values – but, do they share equally? • Consider HCl: H = EN of 2.1 and Cl = EN of 3.0 • Electrons are pulled, as in a tug-of-war, between the – the bond is polar atoms nuclei – the acquires a slight negative charge, and the – In equal sharing (such as diatomic molecules), the hydrogen a slight positive charge bond that results is called a nonpolar covalent bond • Only partial charges, much less than a • When two different atoms bond covalently, there is an true 1+ or 1- as in ionic bond d+ d- unequal sharing; electronegativity is the ability of an • Written as: H Cl atom in a molecule to attract shared electrons to itself • the + and – signs (with the lower case – the more EN atom will have a stronger attraction, and delta: d + and d - ) denote partial charges. will acquire a slightly negative charge • Can also be shown: – called a polar covalent bond, or simply polar bond. H Cl – the arrow points to the more EN atom

Attractions Between Molecules Polar molecules • They are what make solid and liquid molecular compounds • The effect of polar bonds on the polarity of the entire possible. molecule depends on the molecule shape • Intermolecular attractions are weaker than either ionic or covalent – water has two polar bonds and a bent shape; the highly bonds. electronegative oxygen pulls the e- away from H = very • There are two major types: van der Waals forces (of which there polar! are two) and hydrogen bonding. 1. Dispersion forces (van der Waals ) OR OR • the weakest of all, caused by motion of e- • increases as # e- increases • halogens start as gases; Br is liquid; I is solid – all in Group 17 2. (van der Waals force) occurs when polar molecules • When polar molecules are placed between oppositely are attracted to each other, like in water. charged plates, they tend to become oriented with • + region of one molecule attracts the - region of another respect to the positive and negative plates. d+ d- molecule H F • occur when polar molecules are attracted to each other • slightly stronger than dispersion forces • opposites attract, but not completely hooked like in ionic solids

Attractions Between Molecules Hydrogen Bonding 3. Hydrogen bonding is the attractive force caused by (Shown in water) hydrogen bonded to N, O, F, or Cl • N, O, F, and Cl are very electronegative, so this is a very strong dipole. • And, the hydrogen shares with the lone pair in the molecule next to it. + - • This is the strongest of the intermolecular forces. d d H O Order of Intermolecular Attraction Strengths + This hydrogen is bonded – Weakest are the dispersion forces H d Hydrogen bonding covalently to: 1) the highly – A little stronger are the dipole negative oxygen, and 2) a allows H2O to be a – Strongest is the hydrogen bonding liquid at room temp nearby unshared pair. – All of these are weaker than ionic bonds and pressure.

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