Ionic and Covalent Bonding

Chemical Bonds: Ionic and Covalent

Introduction to Chemical Bonding

• Compounds are formed through tightly bonded chemicals, unlike mixtures which can be easily separated.
• Two primary models for chemical bonding:
  • Ionic bonding.
  • Covalent (or molecular) bonding.
• Covalent bonding is more prevalent.
• While distinct, ionic and covalent bonding share some fundamental similarities.

Ionic Bonding and Valence Electrons

• Ionic bonding relies on the outermost electrons, known as valence electrons.
• When determining valence electrons, focus on the highest energy level, regardless of filling order (e.g., for $4s^2 3d^{10} 4p^2$, the 4th energy level is of interest).
• Valence electrons dictate chemical properties.

Lewis Dot Diagrams

• Professor Lewis introduced a method to visualize s and p electrons, especially for representative elements (Groups 1A to 8A).
• The element symbol is used as the base, with dots representing valence electrons.

Constructing Lewis Diagrams

• Example: Lithium (Li) with electron configuration $1s^2 2s^1$ has one valence electron in the 2nd energy level, so its Lewis diagram is Li•.
• Example: Sodium (Na) with electron configuration $1s^2 2s^2 2p^6 3s^1$ has one valence electron in the 3rd energy level, so its Lewis diagram is Na•.
• Elements in the same group (column) have similar Lewis diagrams (e.g., all Group 1A elements have one dot).

Periodic Table and Valence Electrons

• The A-group number corresponds to the number of valence electrons (e.g., Group 2A elements like beryllium have two valence electrons).
• Boron (Group 3A) has three valence electrons.
• Fluorine, chlorine, bromine, and iodine (Group 7A) have seven valence electrons.
• Noble gases (Group 8A) have eight valence electrons (except for helium, which has two).

Significance of Valence Electrons

• Valence electrons are crucial because they mediate interactions between atoms.
• Inner electrons are shielded and do not participate in bonding.

Octet Rule and Isoelectronic Configurations

Achieving Noble Gas Configuration

• Atoms tend to gain, lose, or share electrons to achieve a stable noble gas configuration (8 valence electrons, or 2 for elements aiming for helium's configuration).
• This is known as the octet rule.

Examples

• Fluorine (7 valence electrons) needs one more electron to achieve an octet.
• Sodium (1 valence electron) can lose one electron to achieve the electron configuration of neon.

Electron Transfer and Ion Formation

• When sodium and fluorine react, sodium donates its valence electron to fluorine.
• Sodium becomes a sodium ion ($Na^+$, 11 protons, 10 electrons).
• Fluorine becomes a fluoride ion ($F^-$, 9 protons, 10 electrons).
• Ions are charged atoms formed by gaining or losing electrons.
• Lewis diagrams of ions are enclosed in brackets with the charge on the outside.

Terminology

• Cations: Positively charged ions (e.g., $Na^+$, the 't' in cation resembles a plus sign).
• Anions: Negatively charged ions (e.g., $F^-$).
• Isoelectronic: Having the same electron configuration (e.g., $Na^+$, $F^-$, and Ne are isoelectronic).

Duet Rule

• Lithium (Li) loses its one valence electron to become isoelectronic with helium, following the duet rule.
• Lewis diagrams for positive ions often omit the valence electrons, while negative ions show the complete octet.

Formation of Ionic Compounds

• Ions of opposite charges attract, leading to the formation of ionic compounds.
• Example: Sodium ions ($Na^+$) and fluoride ions ($F^-$) attract to form sodium fluoride (NaF).
• Ionic compounds consist of a lattice of alternating positive and negative ions in a three-dimensional arrangement.
• The chemical formula represents the ratio of ions (e.g., NaF is a 1:1 ratio).

Examples with Polyatomic Ions

• Magnesium (Mg) has two valence electrons and chlorine (Cl) has seven.

• Magnesium donates one electron to each of two chlorine atoms, forming $Mg^{2+}$ and two $Cl^-$ ions.

• The resulting compound is magnesium chloride ($MgCl_2$), with a 1:2 ratio.

• Aluminum (Al) has three valence electrons and oxygen (O) has six valence electrons.

• Two aluminum atoms donate electrons to three oxygen atoms, forming $Al^{3+}$ and $O^{2-}$ ions.

• The resulting compound is aluminum oxide ($Al<em>2O</em>3$), with a 2:3 ratio.

Writing Formulas for Ionic Compounds

• The metal (cation) is written first, followed by the nonmetal (anion).

Exceptions and Considerations

• Carbon (C) typically does not form ionic compounds.
• Elements can lose one, two, or three electrons but rarely more due to increasing ionization energy.
• Elements in Groups 5A, 6A, and 7A typically gain electrons but do not lose them.
• Germanium, tin, and lead can form ions more readily than carbon and silicon due to weaker hold on outer electrons.

Chapter 3 Review: Atomic Structure

• Atomic structure: particle size, charge, location. Tell me about the nucleus.
• Electrons are in the orbital. And that would get to that at the end, but they're out there. And what is most of the whole volume of this big electron?
• Photons and neutrons: No. That's the mass. 4,000 times small. It's eighteen thirty, so you don't need to know that.
• Magnetism: You won't need to know about it. But electron when we get to the orbital diagrams, when you see are separate electrons parallel spins.
• Electrons: Negative. And in the neutral atom, tell me the relationship between protons and electrons. They're the same number.
• In the neutral atom. Because chapter four, we just disturb that neutral atom and we gain and lose electrons. So that's no longer true. But in this chapter, neutral atom undisturbed, say, if it's sodium 11, it's gonna have 11 electrons. 11 protons, 11 electrons.
• Nucleus: Tiny in the center, high dense packing. It was a science fiction movie. I don't even remember the name of it. It's kinda silly. But they used I liked it because they used the idea that atoms are mostly empty space.
• Isotopes: Isotopes are the same element. So if I tell you there's a fluorine nine twenty and there's a fluorine nine twenty one. How do they differ? Same number of Protons. Different number of Neutrons. And how many neutrons in this guy? 11. This is look at your periodic table. Find fluorine. What's the number? 9.
• Whole number is the number of the atomic number is the number of Proton. Protons. Every element has an atomic number. Hydrogen is one. What does that represent? 21. That's a rare isotope made in atomic labs. What's the number of protons? One. One. Atomic number. Mass number, one proton, and how many neutrons? Two. Two.
• Questions about the atom, the subatomic particles, the mass, the volume, and there'll be something on one of these, how many electrons, protons, neutrons, whatever.
• Four families: Not counting hydrogen, group one. Noble gas. Noble gases. So you know a couple other questions that will be on the test. No surprises. These are things you should know. And how would you tell which is a metal and which is a nonmetal? How about indium? Metal or nonmetal? Metal metal metal. Because it is to the left of the staircase. And if it's left and right, and you know you're left and right. And and, you know, iron is a metal. It's on the same side as indium, the elements. You don't have to know them. You just know everything to the left. And to the right, nonmetals. And, really, we didn't do much with that. And then we just did it today.

Electron configuration

• Copper: 29.
• 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d9. One of these electrons can jump here because the three d 10 has added stability. Just this. You don't need to know this, but it's kinda cool. Now how many electron dots would copper have? One. And if he loses that, he's a copper Copper. Plus one. So copper, it turns out, has two different ions.
• Copper has two different forms. It'd be hard to name. We'll have to learn a special way of naming. That's who we are. But and that's true of a lot of those elements in there in the deep lab. But all you need would need to know is how to do this.
• If you got 35, or four, you're not gonna have any