Chemistry: Week 2 - Ionic and Covalent Compounds

Ionic vs. Covalent Compounds

• Building on the understanding that ionic and covalent compounds differ.
• Covalent compounds consist only of nonmetals (including hydrogen).
• Ionic compounds contain a metal and a nonmetal element.

Ionic Compounds: Electrolytes and Salts

• Ionic compounds are also known as electrolytes or salts:
  • They conduct electricity in solution, crucial for bodily functions and batteries.
  • Needed to balance fluids, maintain blood pH, and carry electrical signals in the body.
  • Examples include sodium ($Na^+$), chloride ($Cl^-$), magnesium ($Mg^{2+}$), calcium ($Ca^{2+}$), and potassium ($K^+$).
• Ethanol (covalent) is a poor conductor of electricity, unlike sodium chloride (ionic).
• Ionic compounds separate into ions in solution, enabling electrical conductivity.
• Formed through electrostatic attraction: cations (+) and anions (-) attract strongly; strongest bond type.
• High melting points: requires much more heat to melt salt than sugar, due to stronger bonds.
• Chemical Symbols:
  • Number of atoms in the formula is shown in the bottom right corner (subscript).
  • Charge is in the upper right corner, including +/- and the number value.
  • For a charge of one, the number is often omitted (e.g., $Na^+$).

Formation of Ionic Compounds Involving Valence Electrons

• Valence electrons are involved in bond formation.

  • Sodium (Na) with one valence electron aims for a full octet (8 electrons) like noble gasses.
  • Chlorine (Cl) with seven valence electrons also seeks a full octet.

• Sodium transfers its electron to chlorine:

  • Sodium achieves a full octet by losing one electron, becoming positively charged ($Na^+$, a cation).
  • Chlorine gains an electron to complete its octet, becoming negatively charged ($Cl^-$, an anion).
  • The resulting positive and negative charges create a strong bond.

• In chemical formulas, the cation is always listed first, followed by the anion.

• Charges must cancel out for the compound to be neutral.

  • Example: Oxygen (6 valence electrons) + Lithium (1 valence electron).
  • Lithium aims for a duplet (2 electrons) like Helium.
  • Oxygen seeks an octet (8 electrons) like Neon.
  • Two lithium atoms each donate one electron to oxygen, resulting in $Li_2O$ (lithium oxide).
  • Charge balance: +2 (from two Li+) and -2 (from $O^{2-}$), resulting in a net zero charge.

Naming Ionic Compounds

• When charges are equal and opposite (e.g., Na+ and Cl-), compounds combine in a 1:1 ratio (NaCl).
• Name the cation first, then the anion with an "ide" ending (e.g., sodium chloride), no need for prefixes.
• Example: Lithium Bromide (LiBr); Calcium Bromide (CaBr2).
• Naming Practice:
  • $K_2O$: Potassium Oxide
  • $Al<em>2S</em>3$: Aluminum Sulfide
  • $Ca<em>3N</em>2$: Calcium Nitride
  • LiF: Lithium Fluoride

Determining Formulas for Ionic Compounds: The Charge Flip Shortcut

• When charges are not equal and opposite, use the "flipping the charges" shortcut.
• Calcium ($Ca^{2+}$) and Chloride ($Cl^−$) have unequal charges (+2 ≠ -1).
• The absolute value of the charge of the cation becomes the number of anions, and vice versa.
• The charge and sign is discarded, only the number is used for the subscript to indicate number of atoms since Negative atoms are not possible
• Calcium chloride becomes $CaCl_2$ (calcium: +2, chloride: -1 x 2 = -2, total charge = 0).
• Detailed Example: Determining Aluminum Oxide Formula.
  • Aluminum ion ($Al^{3+}$) and Oxide ion ($O^{2−}$) charges need to balance.
  • Adding more negative charges or more positive charges until they are even.
  • The charge of the oxide ion becomes the number of aluminum ions, and vice versa.
  • The final formula is $Al<em>2O</em>3$ (Aluminum Oxide).

Applying the Shortcut

• Steps:
  • Write out the ions and their charges.
  • Flip the charges to determine the number of each ion in the formula.
• Practice:
  • Potassium Fluoride: K+ and F− → KF
  • Calcium Phosphide: $Ca^{2+}$ and $P^{3−}$ → $Ca<em>3P</em>2$
  • Sodium Sulfide: Na+(+1) and S(2−) → $Na_2S$
  • Aluminum Nitride: $Al^{3+}$ and $N^{3−}$ → AlN

Polyatomic Ions

• Ions containing multiple atoms bound together that act as a single unit with a specific formula and charge.
• Most polyatomic ion names end in "-ate" or "-ite," not "-ide" (except hydroxide).
• Examples: Calcium carbonate, magnesium hydroxide.
• Polyatomic ions are provided on the periodic table during exams; recognition is essential.
• Follow the same rules as before, charges still need to be balanced and neutral.
• Parentheses are used when more than one polyatomic ion is present in a formula, the subscript distributes to every atom within the parentheses.
• Name the polyatomic ion as a single ion.
• Example: Calcium Nitrate $Ca(NO<em>3)</em>2$
• Formulas are kept consistent to aid recognition of polyatomic ions, like nitrate ($NO_3$).
• Naming Practice:
  • $Na<em>3PO</em>4$: Sodium Phosphate
  • $NH_4Cl$: Ammonium Chloride
  • $Mg(NO<em>2)</em>2$: Magnesium Nitrite

Properties of Ionic Compounds

• High melting points.
• Crystalline structure: contributes to stability.
• Dissolve well in water; separable into ions (electrolytes).
• Hard but brittle texture.

Summary of Ionic Compounds

• Formed between a metal and a nonmetal.
• Full transfer of electrons is required, one atom must receive and the other must give away their electrons.
• Held together by strong electrostatic interactions and have to be neutral.
• Name the ions, cation first, anion with "-ide" ending.
• Polyatomic ions have special names and charges.
• High melting points, dissolve well in water, crystalline structure.
• Do not use number prefixes.

Covalent Bonding

• Involves only nonmetals.
• Electrons are shared instead of transferred and the melting points are low because bonds are weaker.
• Some are electrolytes if they dissolve in water; some are not.
• Multiple possible formulas between the same two nonmetals which can be tricky.
• Example: Carbon and oxygen can form carbon monoxide (poisonous) or carbon dioxide.

Covalent Bonds and Water Molecules

• Covalent bonds involve sharing electrons, only nonmetals.
• Nonmetals include hydrogen, and groups 4A, 5A, 6A, 7A, and 8A on the periodic table.
• Water ($H_2O$) Water molecules consist of two hydrogens and one oxygen.
• The oxygen atom has eight electrons in its outermost valence shell because it’s sharing electrons with the hydrogen. Hydrogen shares electrons with oxygen needing a duplet to be stable.
• Bonds are weaker than ionic bonds because there is no full charge.

Number Prefixes in Covalent Compound Names

• Prefixes are used to distinguish compounds with different numbers of the same elements because of multiple possibilities between two nonmetals.
• Unlike ionic compounds which determine based on charge, you must memorize the table for number prefixes.
• List of Numbers: 1-Mono, 2-Di, 3-Tri, 4-Tetra, 5-Penta, 6-Hexa, 7-Hepta, 8-Octa, 9-Nona, 10-Deca.
• In the first atom of a covalent formula, if it’s a number 1, do not have to say mono.
• If there are four carbons and eight hydrogens, then it is named tetracarbon octahydrate.

Naming Covalent Compounds: More Practice

• Naming Practice and Tips:
  • $NHO_3$: Nitrogen Trihydride
  • CO: Carbon Monoxide
  • $CO_2$: Carbon Dioxide
  • $S<em>2F</em>5$: Disulfur Pentafluoride
• In covalent compounds remember to suffix the second part of the name with “-ide”.
• If you see number prefixes then it’s definitely a covalent bond.

Chemical Formulas for Covalent Compounds

• If given the name, determining the chemical formula is very simple. All that is needed is to write down how many there are of something.
• Examples:
  • Carbon Tetrahydride = $CH_4$
  • Dihydrogen monosulfide = $H_2S$

Lewis Dot Structures

• They show the number of valence electrons around the element symbol.
• One dot is one electron, four electrons per side, so do not draw the second dot until each side has one
• Paired dots are not used for bonds but unpaired dots can be used to form covalent bonds.
• The more unpaired dots, the more bonds it can form. You can also determine from the lewis dot structure which electrons will be shared.
• When it comes to bonds, covalent is sharing and ionic is taking.

Lewis Dot Examples

• Helium: Has two valence electrons, neither one is paired up. It has its duplet!
• Carbon: Has four valence electrons, so it can form UP TO four bonds because it wants an octet, so it has four unpaired single dots.
• Neon: Has eight valence electrons because it is in group 8A and octet is already full, so it cannot form bonds with pairs of electrons.

Covalent Bonds: Sharing, Single, Double, and Triple Bonds

• Covalent bonds occur between two nonmetals and only the single, unpaired valence electrons are going to be used in the bonds.
• Fluorine and Chlorine each have one unpaired electron, so they can bond together, that’s where you get two electrons to pair and make a covalent bond.
• Covalent bonds can be single, double, or triple. As you add more bonds, the bonds get stronger and also get shorter.
• Single bonds can occur between two unpaired electrons, double bonds require four unpaired electrons, triple bonds involves six unpaired electrons.
• Number of unpaired electrons = Number of bonds it can form. You can tell all of this from the lewis dot structure.

Covalent Bonds Summary

• Bonds occur between nonmetals, involve the sharing of electrons. When naming, the first nonmetal retains its name, the second gets ized. In naming, add those number prefixes. Remember that LEWIS DOT Structures determine bonds based on paired or unpaired electrons!

Molecular Compounds: Shape and Polarity

• Shapes help determine the overall structure which will affect our health, especially with medicines in health care. There might be medicines that mimic poison to recover your body! Important to know the shape of compounds because of different interactions in the world.
• Polarity: Like Dissolves Like! Water dissolves Water because it is very polar, however oil does not dissolve in water because oil is nonpolar. Polarity is super important for cell structure, and understanding shapes!

Shapes

• Overall Shapes are determined by the Lewis Dot Structure
• Long Way to Determining Shape:
  • How many unpaired valence electrons do they have? Oxygen has six, Hydrogen has 1.
  • The central atom can form the most bonds, or the one there is only one of in the formula because of bonds formed. Oxygen is the central oxygen in $H_2O$!
  • Connect the dots in order to get the shape! Not used in the bonds are two pairs that were paired up already!
• Short Cut way – There is a handy chart to find the short cuts
  • Determine the central atom in the formula FIRST! So in the example $H^2O$, the atom is oxygen! The number of atoms around the atoms! This is also for the handy chart!
• In $CH_4$, Carbon (C) is in the middle because it can form the most bonds of atoms. If you have four atoms attached to it, you have four atoms tetrahedron! Because it does not have nonbonded electrons!
• Extra tip, if there are extra dots, that means it is either a double or triple bond!

Molecular Shape ShortCut Summary

• Overall Summary:
  • Determine the central atom, chemical formula shows the number of atoms around it
  • To determine what it means, let’s determine if has a paired electron pair structure. Example: Bent!

Polar Bonds

• Polar Bounds are how they share with each other
• Are not found in ionics.
• Polar vs. Nonpolar - All about how well they share with each other.
• Polar covalent compounds involve two different nonmetals, and they are gonna have a difference in negativity between the atoms and the electrons are not shared equally.
  • Results are slightly negative side and are only gonna solve in solution.
• Nonpolar covalent bonds are not going to have a significant difference -- the sharing will be awesome. Either involves one type nonmetal , so don't share equal and no balls! Nonpolar things are only gonna dissolve in nonpolar connections.
  *Important! Think back to today and determine if nonpolar or polar/covalent bond?
  Polar covalent bond means you have two different nonmetals!
  Also note:
  electronegativity increases to the right and to the top of the periodic chart.