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Ch 6 Chemical Bonding What You Should Learn in This Section (Objectives): Introduction to Chemical Bonding Types of Chemical

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Ch 6 Chemical Bonding What You Should Learn in This Section (Objectives): Introduction to Chemical Bonding Types of Chemical Ch 6 Chemical Bonding What you should learn in this section (objectives): Define chemical bond Explain why most atoms form chemical bonds Describe ionic and covalent bonding Explain why most chemical bonding is neither purely ionic or purely covalent Classify bonding type according to electronegativity differences. Introduction to Chemical Bonding There are very few atoms that exist as individual particles in nature. Most atoms are bonded to other atoms to form compounds. A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. One reason atoms bond is to decrease their amount of potential energy. When atoms exist by themselves they have relatively high potential energy. Nature favors arrangements that have minimum potential energy. When atoms bond it decreases the amount of potential energy and creates a more stable arrangement of matter. It takes less energy to bond atoms together than to break the bonds between atoms. Bond energy is the amount of energy it takes to break a chemical bond. Types of Chemical Bonds Ionic bonding- chemical bonding that results from the electrical attraction between cations and anions. In purely ionic bonding atoms completely give up electrons to other atoms. Ionic bonding generally involves metals and nonmetals Covalent bonding- the sharing of electron pairs between two atoms. In a pure covalent bond the electrons are shared by the two bonded atoms. Covalent bonding generally involves two nonmetals. A nonpolar-covalent bond is when the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charges. When the distribution of charge is uneven we call this polar. Polar covalent bonds occur when the bonded atoms have an unequal attraction for the shared electrons. Ionic or Covalent? Most chemical bonds are somewhere in between purely ionic and purely covalent. We use the difference in electronegativity values to determine the type of bond that is formed. Remember from Ch 5 that electronegativity is an atom’s ability to attract electrons to its self. We use the following image to determine the bond type. Ionic- electronegativity difference is greater than 1.67 Polar covalent- electronegativity difference is less than 1.67 Nonpolar covalent-electronegativity Problem A difference is less than 0.4 Use electronegativity differences to classify bonding between sulfur and the following elements: hydrogen, cesium, and chlorine. In each pair, which atom will be more negative? Covalent Bonding and Molecular Compounds What you should learn in this section (objectives): Define molecule and molecular formula Explain the relationships among potential energy, distance between approaching atoms, bond length and bond energy. State the octet rule List the six basic steps used in writing Lewis structures Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both. Explain why scientists use resonance structures to represent some molecules. A molecule is a neutral group of atoms that are held together by covalent bonds. They can exist as two or more of the same elements bonded together or two or more different elements bonded together. A chemical compound whose simplest units are molecules is called a molecular compound. A chemical formula shows the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. Remember sub means below. A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound. Formation of a Covalent Bond Nature favors chemical bonding because most atoms have lower potential energy when they are bonded to other atoms than they have as they are independent particles. When 2 hydrogen atoms approach each other 2 “bad” things happen: electron/electron repulsion and proton/proton repulsion. One “good” thing that happens: proton/electron attraction. When the attractive forces offset the repulsive forces, the energy of the tow atoms decreases and a bond is formed. Remember, nature is always striving for a lower energy state. too CLOSE too FAR just RIGHT Bond length is the distance between the two nuclei where the energy is minimal between the two nuclei. In other words, it is the average distance between two bonded atoms. When bonds form individual atoms release energy as they change from isolated individual atoms to molecules. Bond energy is the amount of energy that is required to break the bond. The units for bond energy are usually kj/mol (kilojoule per mole). Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level (outer most ring of the atom). • When two atoms form a covalent bond, their shared electrons form overlapping orbitals. • This achieves a noble-gas configuration. • The bonding of two hydrogen atoms allows each atom to have the stable electron configuration of helium, 1s2. Exceptions to the Octet Rule Fewer than 8: H at most only 2 electrons (one bond),BeH2, only 4 valence electrons around Be (only 2 bonds), Boron compounds only 6 valence electrons (three bonds) Expanded valence (more than 8): can only happen if the central element had d-orbitals which means it is from the 3rd period or greater and can thus be surrounded by more than four valence pairs in certain compounds. The number of bonds depends on the balance between the ability of the nucleus to attract electrons and the repulsion between the pairs. Some of the more elements are fluorine, oxygen, chlorine and noble gases. Electron-Dot Notation • To keep track of valence electrons, it is helpful to use electron-dot notation. • Electron-dot notation is an electron- configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol. The inner-shell electrons are not shown. Lewis Structures H:H An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom. Lewis Structures are formulas in which atomic symbols represent nuclei and inner-shell electrons, dot- pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons. A structural formula indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule. Example F-F and H-Cl. Single bonds (sigma bonds) are a covalent bond in which one pair of electrons is shared between two atoms. They are represented by two dots (electrons) or one dash. These are the longest bonds, but also the weakest Double bonds (pi bonds) are covalent bonds in which two pairs of electrons are shared between two atoms. They are represented by four dots (electrons) or 2 dashes = Example C=C Triple bonds (pi bonds) are covalent bonds in which three pairs of electrons are shared between two atoms. They are represented by six dots (electrons) or 3 dashes. Example These are shortest, but also the strongest. Carbon, nitrogen, oxygen, phosphorous, and sulfur are the most common elements that form multiple bonds. Drawing Lewis Structures 1. H is always a terminal atom. ALWAYS connected to only one other atom. 2. Lowest electronegativity is the central atom in a molecule. 3. Find the total number of valence electrons by adding up group numbers of the elements. For ions add electrons for negative charges and subtract electrons for positive charges. 4. Place one pair of electrons (sigma bond) between each pair of bonded atoms. 5. Subtract from the total number of bonds you just used. 6. Place lone pairs about each terminal atom (except H) to satisfy the octet rule. Left over pairs are assigned to the central atom. 7. If the central atom is not yet surrounded by four electron pairs, convert on or more terminal atom lone pairs to a double or triple bond ( pi bonds). Only C, N, O, P, and S can form multiple bonds (pi bonds) Resonance Structures Resonance refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure. Ozone (O3) exists as an average of these two images, so it must be shown both ways. Ionic Bonding and Ionic Compounds What you should learn in this section (objectives): Compare and contrast a chemical formula for a molecular compound with one for an ionic compound. Discuss the arrangements of ions in crystals Define lattice energy and explain its significance List and compare the distinctive properties of ionic and molecular compounds Write the Lewis structure for a polyatomic ion given the identity of the atoms combined and other appropriate information. Formation of an Ionic Compound Ionic compounds are composed of positive (cation) and negative (anion) ions that are combined so that the numbers of positive and negative charges are equal Ionic Bond - Completely transfer electrons. Positive charge – cation – lost electrons to the anion. Negative charge – anion – gained electrons from the cation. Positive charge must equal and, therefore, cancel the negative charge. Example: Sodium Chloride – sodium wants to lose one electron to become stable and chlorine wants to gain one electron to become stable. Formula unit – a chemical formula of the smallest sample of an ionic compound. Ionic Character Ionic compounds have the greatest ionic character with full on charged ions. The further the ions are apart in electronegativity, the more the ionic character. Molecular compounds have very low electronegativity. The closer the ions are in electronegativity, the less the ionic character. Characteristics of Ionic Bonding Ionic compounds are crystalline solids at room temperature. They are arranged in repeating three-dimensional pattern called a crystal lattice.
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