Structure Properties of Matter Lesson 4 Chemical Bonding

Structure & Properties of Matter Lesson # 4: Chemical Bonding

Definitions Chemical Bond – the electrical attraction that holds atoms or ions together in a molecular element or in an ionic or molecular compound. • Molecular element – bonding between two or more of the same atom • Molecular compound – bonding between two or more different atoms • Ionic compound – pure substance composed of two or more ions combined in a fixed ratio. •

Ionic Bonding • Chemical bond between oppositely charged ions. They form when an atom that loses electrons relatively easily reacts with an atom that gains electrons relatively easily. • The transfer of electrons occurs because of one atom has a much higher electronegativity than the other. • Typically, an ionic compound results when a metal element reacts with anon-metal element to form positively charged and negatively charged ions that are held together by electrostatic attraction.

Crystal Lattice – This attraction creates a crystal lattice structure. – This configuration is more stable than if the ions were separated into individual formula units. – It is arranged so that the attractions between opposite charges are maximized and repulsions between like charges are minimized.

Ionic Bonding • When atoms form ionic compounds, they become isoelectronic with a noble gas (it has the same electron configuration as the closest noble gas). – Example: Calcium: [Ar] 4 s 2 and Oxygen: [He] 2 s 2 2 p 4 • Calcium needs to lose its 4 s shell to have a stable n=3 orbital (lose 2 electrons), and oxygen needs to gain two to fill its n=2 orbital. • Calcium transfers its two electrons to oxygen, resulting in calcium with the same electron configuration as argon, and oxygen as neon, both stable noble gases.

Molecular Elements • When hydrogen atoms come close to each other, the attraction between its one electron and the proton in the nucleus in an adjacent hydrogen atom creates a force that pulls both atoms toward each other. • At a certain distance, there exists a perfect balance between attractive and repulsive forces creating a molecule, H 2, that is more stable than when the atoms are apart. – This is an example of a molecular element, and in this case, a diatomic molecule.

Covalent Bonding • The hydrogen atoms are held together by a covalent bond, which forms when atoms share electrons. • Each set of shared electrons are called electron bonding pairs. • Covalent bonds typically form between two or more non-metal atoms.

Lewis Structures • In 1916, Gilbert Lewis provided a theory of bonding. His postulates included: – Atoms are ions are stable if they have a full valence shell of electrons – Electrons are most stable when they are paired – Atoms form chemical bonds to achieve a full valence shell of electrons – A full valence shell of electrons may be achieved by an exchange of electrons between a metal and non-metals atoms. – The sharing of electrons results in a covalent bond.

Lewis Structures • A Lewis structure shows the arrangement of electrons and bonds in a molecular or polyatomic ion. • The first row elements (H, He) on the period table follow the duet rule – need to have two electrons in their valence shell to be stable. • Elements in period 2 follow the octet rule – stating that many atoms form the most stable substances when they are isoelectronic with neon, which has 8 electrons in its valence shell.

Example: Fluorine (F 2) – Even though each fluorine atom shares 1 of its electrons with the other, each also has 3 pairs of electrons that are not involved in bonding – they are called lone electron pairs.

1. 2. 3. 4. Rules for Drawing Lewis Structures Identify the central atom – the element with the highest bonding capacity. Draw the structure and arrange the other atoms around it. Add up the number of valence electrons for each atom so you know the total number of dots to add into your structure. If the structure is a polyatomic ion, add or subtract the appropriate number of electrons. Add in one bond with 2 dots, and place lone pairs around the outer atoms as necessary (following the duet and octet rule). See how many more electrons are necessary to add in, and add multiple bonds to the central atom if necessary. Don’t forget to use square brackets for ions.

Example – Methanal (CH 2 O)

Example – Nitrate Ion (NO 3 -)

Exceptions to the Octet Rule – Hydrogen, Carbon, Nitrogen, Oxygen, and all the halogens always follow the octet rule – no exceptions. – Other atoms can break the octet rule, forming molecules with fewer or more than 8 electrons surrounding them. – Beryllium, for example, is happen with four electrons in its outer shell. Boron and aluminum typically form compounds with only 6 valence electrons. – These compounds are typically quite reactive as they are electron deficient, and will react with electron-rich compounds like water and ammonia to form a compound that does obey the octet rule.

Example – Boron Trifluoride

Exceeding the Octet Rule – Some atoms can exceed the octet rule. This happens because they have a nearby vacant d orbital to exceed their octet. – This will only happen in period 3 or higher, as period 1 and 2 elements do not have d orbitals. – The extra electrons beyond 8 will take residence in the empty d orbital. Example – sulfur has electrons in its 3 s and 3 p orbital. They extra electrons can fill out the 3 d orbital at a slightly higher energy level.

Example: Sulfur Hexafluoride

Coordinate Covalent Bonding • Most covalent bonds are formed from one electron pairing with another from a different atom. • When both electrons come from one atom, it is called a coordinate covalent bond. • For example, the ammonium ion (NH 4+) forms when a hydrogen ion interacts with the lone pair on the nitrogen in an ammonia molecule. The hydrogen ion’s positive charge is attracted to the lone pair, and forms a coordinate covalent bond.

Example: Ammonia + Hydrogen Ion

Resonance Structures • Sometimes a Lewis structure that meets all criteria is not supported by experimental observations. – For example, the structure for ozone, O 3 (g), seems it should have one single and one double bond when the Lewis structure is drawn. – However, experimental measurements of bond lengths indicate that the bonds between the oxygen atoms are actually identical and have properties that are somewhere between a single bond a double bond.

Ozone – To communicate the bonding in O 3 more accurately, chemistry draw two Lewis structures: – The double headed arrow indicates that the actual structure is between these two structures, not that the structure is changing back and forth. Each of these Lewis structures is called a resonance structure, which is one of two or more structures that differ only in the position of their bonding and lone pairs.

Resonance Hybrid The real molecule are actually a combination of the two structures (almost like 1. 5 bonds), and is called a resonance hybrid.

Example: Carbonate Ion (CO 3 -) 2