Chapter 2: Bonding (C4041645)

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Getting started

In chemistry, bonding is the connection of two atoms.

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Frequently asked questions
What is bonding?
Where two atoms are adjoined.

1 Ionic bond

Ionic bond is a bond between two oppositely charged ions. In the bond, one element donates electron(s) to another element. Ions are created when an atom or molecule have unbalanced positive and negative charges. Cations are positive ions, and anions are negative ions. Ionic compounds are unique as they conduct electricity when in molten or solution, but not in solid. They tend to have a high melting point, and be soluble in water.

Frequently asked questions
What is an ionic bond?
It's a bond between 2 oppositely charged ions.

What's an ion again?
It's an atom or molecule that's charged. So in an ionic bond, one is positively charged, and the other is negatively charged.

What is a cation?
A positively charged ion.

And an anion is a negatively charged ion?
You got it !

Why are ions only conductive under certain circumstances, and not in others?
They have charge, so they can conduct electricity when these ions can move around. They can't move around in a solid, because everything is rigid. But in molten or solution, they can. So that's why ions are only conductive under certain circumstances.

Ions have a high melting point?
Yep, because the ionic bond is a strong bond!

And ions are soluble in water?
You got it! That's because water is slightly charged, and being charged, ions dissolve readily in water.

Formative learning activityMaps to RK2.A
What are ionic bonds, and what is unique about them?

2 Covalent bond

Molecules are repeating groups of two or more atoms, held together by covalent bonds.

Covalent bond are bonds created by the sharing of electron pairs (two electrons, one donated from each atom) between two atoms [as supposed to electron donation in the ionic bond]. Each atom’s protons attracts to both electrons in the electron pair, creating a tug-of-war, thereby bonding the two atoms. Although the nuclei of both atoms repel each other, the nuclei are attracted to the electron pairs. There is a distance between two nuclei which is optimal, known as the bond length. Bond length and bond strength are inversely related, such that as there is decreased bond length, there is increased bond strength, as there is the lowest energy [and therefore most stability] when electrons are closer to their nuclei. Covalent bonds can also be created with two electrons donated from the same atom (rather than donated from each atom), known as a coordinate covalent bond.

Compounds are substances consisting of two or more different elements. They are distinct from molecules, in that they do not have distinct groups.

Lewis [electron] dot structures/formulas illustrate the valence electrons in covalent molecules. Valence is the number of electrons in the outer most shell. Valence can be determined from the periodic table group number, omitting the transition metals from the group count, such that lithium has 1 valence electron, beryllium has 2 valence electrons, boron has 3 valence electrons, carbon has 4 valence electrons, nitrogen has 5 valence electrons, oxygen has 6 valence electrons, fluorine has 7 valence electrons, and neon has 8 valence electrons.

As covalent bonds involve sharing electrons, rather than being like ionic bonds where one atom gains and another loses an electron, both atoms effectively gain an electron. Thus, each atom, like in ionic bonding, can be thought to desire an additional number of electrons to reach its noble gas configuration. The desire to reach nearest noble gas configuration is known as the octet rule. Although octet refers to having 8 electrons in the outer shell, this only applies to the p subshell. From , the s subshell (notably hydrogen) will desire 2 known as a duet, the d subshell will desire 18 known as the 18-electron rule, etc. Thus, fluorine will desire 1 extra electron known as monovalent, oxygen will desire 2 extra electrons known as divalent, nitrogen will desire 3 extra electrons known as trivalent, and carbon will desire 4 extra electrons known as tetravalent. There are exceptions to the octet rule though, including:

  • The statement atoms desire to reach its nearest noble gas configuration, is under-simplified: in fact, bonding is determined by the desire to bond unpaired electrons, known as lone electrons. (Applying the principle as , fluorine has 3 electron pairs, and one lone electron, thereby desiring one bond.) Expansion of the octet, is where some period three elements, bump up electrons from the lower 3s or 3p subshells, into the higher empty 3d subshell. For example, although sulfur can normally form 2 bonds, in [mathjax]\ce{SF6}[/mathjax] it requires 6 bonds, so bumps up an electron from the 3s and 3p subshell, so that instead of being [mathjax]1s^2 2s^2 2p^6 3s^2 3p^4[/mathjax], it becomes [mathjax]1s^2 2s^2 2p^6 3s^1 3p^3 3d^2[/mathjax], noting now that there is 1 lone electron in 3s, 3 lone electrons in 3p, and 2 lone electrons in 3d, meaning there are [mathjax]1+3+2=6[/mathjax] lone electrons, thereby now permitting 6 bonds
  • In line with the restated formula that lone electrons desire to be bound, where a molecule still has non-bonded lone electrons, the molecule is known as a free radical, which are highly reactive. Because bonding causes an electron to be localized, bonding reduces net charge, for example, when [mathjax]O^{2-}[/mathjax] is bound to 2 hydrogens, it becomes [mathjax]\ce{H2O}[/mathjax] with no net charge. Hence, if it is non-bounded, there is a net charge
  • Some atoms have incomplete octets, which do so because they can often react further to complete their octet. For example, boron in [mathjax]BF3[/mathjax] only has 6 electrons in its valence shell
  • Atoms with an odd number of valence electrons, such as nitrogen (with 5 valence electrons)

Formal charge is the charge of an atom in a molecule (not assuming attachments are removed). It is calculated by the number of bonds it can make minus the number of bonds it does make, which is [mathjax]FC=V-N-B[/mathjax], where [mathjax]V[/mathjax] is the number of valence electrons, [mathjax]N[/mathjax] is the number of lone electrons (lone meaning non-bonded), and [mathjax]B[/mathjax] is the number of bonded electrons. For example, in nitrite ion [mathjax]NO2^-[/mathjax], nitrogen is double bonded to one [mathjax]O[/mathjax] (thereby providing 2 bonding electrons), single bonded to an [mathjax]O^-[/mathjax] (thereby providing 1 bonding electron), and has two lone electrons. Nitrogen naturally has 5 valence electrons. Therefore, the formal charge of nitrogen is [mathjax]FC=5-2-3=0[/mathjax]. Another example, in nitrogen dioxide [mathjax]NO_{2} \bullet[/mathjax], which is like nitrite ion in structure, but rather than having 2 lone electrons, has 1 lone electron, thereby since being less negative, therefore positively charges the nitrogen atom. Since there are 2 lone electrons (cf. 1 lone electron in nitrite ion), [mathjax]N=1[/mathjax]. Thus, [mathjax]FC=5-1-3=+1[/mathjax]. Thus, nitrogen dioxide has a formal charge of +1.

In contrast, oxidation number only considers valence. Oxidation number is the charge of an atom in a molecule, if all attachments were removed [and therefore according electron pairs]. The rules of oxidation numbers are:

  • The oxidation number of an element is 0, as there is no charge since there are no attachments
  • When compounded with a non-metal, the oxidation number of hydrogen is +1. When compounded with a metal, the oxidation number of hydrogen is -1
  • Oxygen generally has an oxidation number of -2, but there are exceptions, for example hydrogen peroxide [mathjax]\ce{H2O2}[/mathjax], because hydrogen is assigned with precedence +1, the charge on oxygen will be -2 applying the principles
  • The sign of the oxidative number is equivalent to a hypothetical ionic bonds. For example, in NaCl, chlorine in a hypothetical ionic bond needs to gain an electron, and therefore increases negativity, which is why it is assigned a negative charge when ionically bound. In contrast, sodium needs to lose an electron, and therefore decreases negativity, which is why it is assigned a positive charge when bound. This is why more electronegative, and therefore non-metallic elements, which wish to gain an electron, and therefore increases negativity, are assigned a negative oxidation number. In contrast, less electronegative, and therefore metallic elements, which wish to lose an electron, and therefore lose negativity, are assigned a positive oxidation number
  • The sum of the oxidation numbers of all atoms is its net charge. If there is no net charge, this should be 0. For example, with [mathjax]\ce{CO2}[/mathjax], oxygen is assigned as -2. As there are two oxygen’s, this needs to be accounted for in the equation. Given that the net charge is 0, we can draw that [mathjax]C+2(-2)=0[/mathjax], or reshuffling, [mathjax]C=+4[/mathjax]

Lewis dot structures can be drawn by the process:

  • Determine the number of valence electrons for all atoms
  • Use an electron pair (2 electrons), one from each atom, to form 1 covalent bond between the atoms, such that each atom satisfies the octet rule. A covalent bond is illustrated by a dash
  • Determine the remaining electrons around each atom, and arrange the electrons as lone pairs

For example, in [mathjax]\ce{H2O}[/mathjax], oxygen has a valence of 6, and hydrogen has a valence of 1.Each hydrogen shares one electron with an according electron from oxygen. Although hydrogen no longer has any non-bounded valence electrons, oxygen has 4 (as 2 are bound in the 2 valence bonds with their according hydrogen).

Apart from Lewis dot structures, there are other graphical means of depicting molecular structure, known as structural formulae. There are several formats, including:

  • Skeletal formula, where each carbon is represented by a kink [or end] in a line. Hydrogen’s are not shown, but can be added as required to fulfill the octet rule (in particular, the each carbon makes 4 bonds). Dashed lines represent bonds projecting into the page. Solid lines represent bonds in the plane of the page. Solid wedges represent bonds projecting out of the page
  • Fischer projection, which is a 2D representation of a 3D molecule. Bonds are either horizontal or vertical. Horizontal lines represent bonds projecting out of the page. Vertical lines represent bonds projecting into the page
  • Newman projection, which visualizes a carbon-carbon bond from front-to-back. The closer carbon is represented as a dot, and the further carbon represented as a circle. Newman projects can either be:
    • Eclipsed conformation, where groups are rotated so there is minimal torsion angle. This creates steric hindrance, which is where large groups get close, and thus interfere with each other. Because large groups do not like to be next to each other, there is high conformational energy required to keep the shape
    • Staggered conformation, where groups are rotated so there is maximum torsion angle. Torsion angle is the difference in [rotational] angle between one group of attachments to a carbon, and the group of attachments to another carbon. In contrast with eclipsed conformation, conformational energy is minimum. In particular, antistaggered conformation is when the two largest groups are as far away as possible; and gauche conformation is where the two largest groups are not as far away as possible
  • Ball-and-stick model, which is a 3D representation of a molecule. Atoms are depicted as a ball, and bonds depicted by sticks that connect the balls
  • Space-filling model, which is another 3D representation of a molecule. It is distinct from the ball-and-stick model, in that bonds are not shown. As bonds are not shown, the balls can be larger, and therefore more accurately represent the relative sizes of the constituent atoms, and their distances between each other. The space-filling model most accurately portrays the molecule in reality

Formative learning activityMaps to RK2.B
What are covalent bonds, and what is unique about them?

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Pre-med science (MED5118352)


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Chapter 2: Bonding - General chemistry - Pre-med science - MR. SHUM'S CLASSROOM