How can chemistry be considered the blueprint of matter?
Earth and Atmospheric Sciences
At its simplest, a chemical compound is a substance formed when two or more different chemical elements are chemically bonded together.
Think of elements (like Hydrogen or Oxygen) as individual Lego bricks. A compound is the specific structure you get when you snap those bricks together in a fixed arrangement. Once they are bonded, they lose their individual “personality” and take on entirely new properties.
The Key Ingredients of a Compound
To qualify as a compound, a substance must meet three specific criteria:
- Different Elements: It must contain at least two different types of atoms. For example, O2 (pure oxygen gas) is a molecule, but not a compound, because it only contains one element. H2O (water) is a compound because it has Hydrogen and Oxygen.
- Fixed Ratio: The elements are always present in a specific, unchanging ratio. Water is always two parts hydrogen to one part oxygen. If you change the ratio to H2O2, you no longer have water; you have hydrogen peroxide.
- Chemical Bonds: The atoms are held together by “chemical glue,” usually involving the sharing or exchange of electrons. You can’t separate a compound just by filtering it or boiling it; you need a chemical reaction to break those bonds.
Common Examples
You interact with compounds every second of the day. Here are a few classics:
| Compound | Chemical Formula | What is it? |
| Water | H2O | The universal solvent. |
| Sodium Chloride | NaCl | Basic table salt. |
| Glucose | C6H12O6 | A simple sugar used for energy. |
| Methane | CH4 | The primary component of natural gas. |
| Carbon Dioxide | CO2 | The gas we exhale and plants “breathe.” |
Compounds vs. Mixtures
It’s easy to confuse these two, but there is a big difference in how they “hang out” together:
- Compound: Atoms are bonded. Imagine a cake—once it’s baked, you can’t easily pull the flour back out of the finished slice.
- Mixture: Substances are just mingled. Imagine a bowl of salad. The lettuce and tomatoes are touching, but they haven’t turned into a new substance. You can pick the tomatoes out with a fork.
Why do they form?
Most atoms are “lonely” and unstable on their own. They want to have a full outer shell of electrons (the “Octet Rule”). To achieve this stability, they find partners to swap or share electrons with, resulting in the diverse world of matter we see today.
What are the types of chemical compounds and their formulas?
Wayne Breslyn (Dr. B.)
Chemical compounds are generally categorized by the type of chemical bond that holds their atoms together. Because atoms have different ways of sharing or stealing electrons, they form distinct structures with unique behaviors.
1. Ionic Compounds
Ionic compounds form when one atom (usually a metal) “donates” electrons to another atom (usually a non-metal). This creates charged ions that stick together like powerful magnets.
- Characteristics: High melting points, usually dissolve in water, and conduct electricity when liquid.
- Structure: They form a repeating 3D grid called a crystal lattice.
| Name | Formula | Use |
| Sodium Chloride | NaCl | Table salt |
| Calcium Carbonate | CaCO3 | Chalk and limestone |
| Magnesium Oxide | MgO | Antacids and industrial lining |
| Sodium Bicarbonate | NaHCO3 | Baking soda |
2. Covalent (Molecular) Compounds
These form when two or more non-metals share electrons to stay stable. Instead of a massive lattice, they usually form discrete, individual molecules.
- Characteristics: Lower melting points than ionic compounds; can be solids, liquids, or gases at room temperature.
- Structure: Specific shapes (like “bent” or “linear”) determined by how the electron pairs repel each other.
| Name | Formula | Use |
| Water | H2O | Life’s essential solvent |
| Carbon Dioxide | CO2 | Respiration and carbonation |
| Ammonia | NH3 | Fertilizers and cleaning |
| Glucose | C6H12O6 | Cellular energy/sugar |
3. Organic Compounds
Organic compounds are a special subset of covalent compounds that always contain Carbon (usually bonded to Hydrogen, Oxygen, or Nitrogen). They are the basis of all known life on Earth.
- Note: Not all carbon-containing compounds are organic (like CO2), but almost all organic compounds are built on carbon “skeletons.”
| Name | Formula | Use |
| Methane | CH4 | Natural gas fuel |
| Ethanol | C2H5OH | Alcohol in beverages/fuel |
| Propane | C3H8 | Heating and cooking gas |
| Acetic Acid | CH3COOH | Vinegar |
4. Acids and Bases
These are often categorized separately because of how they react in water.
- Acids usually start with Hydrogen (H) and release H+ ions in water.
- Bases often end with Hydroxide (OH) and accept H+ ions.
| Name | Formula | Type |
| Hydrochloric Acid | HCl | Strong Acid (Stomach acid) |
| Sulfuric Acid | H2SO4 | Strong Acid (Car batteries) |
| Sodium Hydroxide | NaOH | Strong Base (Lye/Drain cleaner) |
How to read a formula
A chemical formula is a shorthand “recipe.”
- Symbols: Tell you which elements are present (C for Carbon, H for Hydrogen).
- Subscripts: Tell you how many atoms of that element are in the compound. For example, in H2SO4, there are 2 Hydrogen atoms, 1 Sulfur atom (the “1” is invisible), and 4 Oxygen atoms.
What is the relationship between the concept of the mole and chemical compounds?
The relationship between the mole and chemical compounds is essentially the bridge between the microscopic world (atoms and molecules) and the macroscopic world (grams and liters).
In chemistry, you can’t count individual molecules—they are far too small. Instead, we use the mole as a “chemist’s dozen.”
1. The Mole as a Counting Unit
Just as a “dozen” always means 12, a mole always means 6.022 * 1023 particles. This is known as Avogadro’s Number.
For a chemical compound, one mole represents that specific number of formula units or molecules.
- 1 mole of Water (H2O): 6.022 * 1023 molecules of H2O.
- 1 mole of Table Salt (NaCl): 6.022 * 1023 units of NaCl.
2. Molar Mass: Turning Formulas into Weight
The chemical formula of a compound tells you exactly how much one mole of that substance weighs. This is called the Molar Mass (measured in g/mol).
To find the molar mass of a compound, you add up the atomic weights of every atom in the formula (found on the Periodic Table).
Example: Calculating Molar Mass for Water (H2O)
- Hydrogen (H): ~1.01 g/mol * 2 atoms = 2.02 g
- Oxygen (O): ~16.00 g/mol * 1 atom = 16.00 g
- Total Molar Mass: 18.02 g/mol
This means if you weigh out exactly 18.02 grams of water in a beaker, you are holding exactly one mole (6.022 * 1023 molecules) of water.
3. The Ratio Within the Compound
The mole concept allows us to look inside a compound to see the ratio of its parts. The subscripts in a chemical formula represent mole ratios.
In 1 mole of Carbon Dioxide (CO2):
- There is 1 mole of Carbon atoms.
- There are 2 moles of Oxygen atoms.
This is vital for scientists because it allows them to calculate exactly how much of a specific element (like the amount of Sodium in salt) is present in a bulk sample.
4. Stoichiometry: The “Recipe” of Reactions
The biggest reason the mole is related to compounds is for chemical reactions. Equations are written in moles, not grams.
Consider this reaction:
2H2 + O2 → 2H2O
This tells us that 2 moles of Hydrogen gas react with 1 mole of Oxygen gas to produce 2 moles of the compound Water. Without the mole, we wouldn’t know the correct “proportions” to mix chemicals without leaving leftovers or causing dangerous imbalances.
Summary Table
| Concept | Relationship to Compound |
| Avogadro’s Number | The number of compound molecules in one mole. |
| Subscripts | The molar ratio of elements within the compound. |
| Molar Mass | The physical weight (in grams) of one mole of the compound. |
| Stoichiometry | Using moles to predict how much compound a reaction will produce. |
What is composition of chemical compounds?
The composition of a chemical compound refers to the specific identity and relative amounts of the elements that make it up. Understanding composition is like knowing both the ingredients and the exact measurements for a recipe.
There are two primary ways chemists describe composition: Qualitative (what is in it) and Quantitative (how much is in it).
1. Elemental Composition (What it’s made of)
At the most basic level, composition is defined by the types of atoms present. This is expressed through the Chemical Formula.
- Atomic Ratio: The subscripts in a formula tell you the ratio of atoms. In Methane (CH4), the composition is always 1 carbon atom for every 4 hydrogen atoms.
- Chemical Identity: Even if two compounds have the same elements, a change in their arrangement or ratio creates a completely different substance (e.g., CO is toxic carbon monoxide, while CO2 is the carbon dioxide we exhale).
2. Percent Composition by Mass
Since different atoms have different weights, chemists often look at the Percentage by Mass of each element. This tells you what fraction of the total “weight” of the compound comes from a specific element.
To calculate the percent composition of an element:
Percent Composition = (Mass of element in 1 mole / Molar mass of the entire compound) * 100
Example: Water (H2O)
Even though there are twice as many Hydrogen atoms as Oxygen atoms, Oxygen is much heavier.
- Mass of Oxygen: ~16.00 g
- Total Mass of H2O: ~18.02 g
- Result: Water is roughly 88.8% Oxygen and only 11.2% Hydrogen by mass.
3. The Law of Definite Proportions
A fundamental rule of chemistry is that a pure compound always contains the exact same elements in the exact same proportions by mass, regardless of where it came from.
Example: A molecule of water taken from the Atlantic Ocean has the exact same chemical composition as a molecule of water vapor on Mars. If the ratio changes, it is no longer water.
4. Empirical vs. Molecular Formulas
Composition can be expressed in two different “depths”:
- Empirical Formula: The simplest, reduced ratio of elements.
- Example: The empirical formula for Glucose is CH2O (1:2:1 ratio).
- Molecular Formula: The actual number of atoms in a single molecule.
- Example: The molecular formula for Glucose is C6H12O6.
Summary Table: Composition of Glucose (C6H12O6)
| Component | Atomic Count | Mass Contribution | Percent of Total |
| Carbon (C) | 6 | 72.06 g | 40.0% |
| Hydrogen (H) | 12 | 12.11 g | 6.7% |
| Oxygen (O) | 6 | 96.00 g | 53.3% |
| Total | 24 Atoms | 180.17 g/mol | 100% |
What are oxidation states and how can it be used to describe chemical compounds?
Think of an oxidation state (or oxidation number) as a “virtual charge” assigned to an atom. It represents the number of electrons an atom gain, loses, or appears to use when joining with other atoms to form a compound.
While atoms in a covalent bond technically “share” electrons, oxidation states pretend the more “greedy” (electronegative) atom takes them all. This helps chemists keep track of electron flow.
1. The Rules of the Game
To determine the oxidation state of an element in a compound, we follow a set of priority rules:
- Free Elements: Any element in its natural, uncombined state has an oxidation state of 0 (e.g., O2, Fe, H2 are all 0).
- Monatomic Ions: The oxidation state equals the charge of the ion (e.g., Na+ is +1, Cl– is -1).
- Oxygen: Usually -2 (except in peroxides like H2O2, where it is -1).
- Hydrogen: Usually +1 when bonded to non-metals and -1 when bonded to metals.
- Fluorine: Always -1 (it is the most electronegative element).
- The Sum Rule: For a neutral compound, the sum of all oxidation states must be 0. For a polyatomic ion, the sum must equal the ion’s charge.
2. Using Oxidation States to Describe Compounds
Oxidation states act as a “ID card” for a compound, telling us about its stability and reactivity.
Identifying the Compound’s Name
In ionic compounds involving transition metals (like Iron or Copper), the oxidation state is included in the name using Roman numerals because those metals can have multiple “moods.”
- FeCl2: Iron(II) Chloride (Iron is +2)
- FeCl3: Iron(III) Chloride (Iron is +3)
Determining Reactivity (Redox)
Oxidation states tell us if a compound is an oxidizer or a reducer.
- Oxidation: An increase in oxidation state (loss of electrons).
- Reduction: A decrease in oxidation state (gain of electrons).
3. Calculating an Unknown State
You can use the “Sum Rule” like a simple algebra equation to find the state of a mystery element in a compound.
Example: Finding Manganese in Potassium Permanganate (KMnO4)
- Potassium (K): Group 1 metal, always +1.
- Oxygen (O): 4 atoms at -2 each = -8.
- Equation: (+1) + Mn + (-8) = 0
- Solve: Mn – 7 = 0 → Mn = +7
In this compound, Manganese is in a very high oxidation state (+7), making it a powerful oxidizing agent.
4. Common Oxidation States Table
| Element | Typical State | Reason |
| Alkali Metals (Na, K) | +1 | Lose 1 electron easily |
| Alkaline Earth (Mg, Ca) | +2 | Lose 2 electrons easily |
| Halogens (Cl, Br) | -1 | Want to gain 1 electron |
| Noble Gases (He, Ne) | 0 | Already stable; rarely form compounds |
Why does it matter?
Without oxidation states, we wouldn’t understand how batteries work, how rust forms, or how our bodies convert glucose into energy. It is the “accounting system” for the most fundamental currency in the universe: the electron.
How are inorganic compounds named?
Naming inorganic compounds follows a systematic set of rules established by IUPAC (International Union of Pure and Applied Chemistry). The method used depends entirely on whether the compound is ionic (metal + non-metal) or covalent (two non-metals).
1. Naming Ionic Compounds
Ionic compounds consist of a positive ion (cation) and a negative ion (anion).
Binary Ionic Compounds (Two Elements)
- Name the Metal first: Use its name exactly as it appears on the periodic table.
- Name the Non-metal second: Change the ending of the element’s name to “-ide.”
- Example: NaCl → Sodium chloride.
- Example: MgBr2 → Magnesium bromide.
Metals with Multiple Charges (Transition Metals)
Many metals (like Iron or Copper) can form ions with different charges. We use Roman numerals in parentheses to specify the oxidation state.+1
- FeCl2 → Iron(II) chloride (Iron has a +2 charge).
- FeCl3 → Iron(III) chloride (Iron has a +3 charge).
2. Naming Covalent (Molecular) Compounds
Since non-metals can combine in multiple ratios (like CO and CO2), we use Greek prefixes to indicate the number of atoms present.
| Number | Prefix | Number | Prefix |
| 1 | Mono- | 6 | Hexa- |
| 2 | Di- | 7 | Hepta- |
| 3 | Tri- | 8 | Octa- |
| 4 | Tetra- | 9 | Nona- |
| 5 | Penta- | 10 | Deca- |
The Rules:
- Apply a prefix to the first element (except if there is only one, omit “mono-“).
- Apply a prefix to the second element and change the ending to “-ide.”
- Example: N2O5 → Dinitrogen pentaoxide.
- Example: CO → Carbon monoxide.
3. Polyatomic Ions
Some compounds contain “clumps” of atoms that stay together as a single charged unit. You must memorize these or use a reference table.
- NaNO3 → Sodium nitrate.
- (NH4)2SO4 → Ammonium sulfate.
- CaCO3 → Calcium carbonate.
4. Naming Acids
Acids are named based on the anion they produce when dissolved in water.
- Binary Acids (H + one element): Use the prefix hydro- and change the ending to -ic acid.
- HCl → Hydrochloric acid.
- Oxyacids (H + polyatomic ion containing Oxygen):
- If the ion ends in -ate, change it to -ic acid. (H2SO4 / Sulfate → Sulfuric acid)
- If the ion ends in -ite, change it to -ous acid. (H2SO3 / Sulfite → Sulfurous acid)
Summary Flowchart for Naming
| If the compound starts with… | It is likely… | Naming Rule |
| A Metal | Ionic | Metal + Non-metal(“-ide”) |
| A Transition Metal | Ionic | Metal(Roman Numeral) + Non-metal(“-ide”) |
| A Non-metal | Covalent | Prefix-Element + Prefix-Element(“-ide”) |
| Hydrogen (H) | Acid | Hydro- prefix or -ic/-ous suffix |