Atoms, bonding, gases, balanced equations, moles, solutions, pH, buffers, equilibrium and an organic and biochemistry preview, with short fill-in-the-blank Python calculators.

Chemistry Fundamentals

## Why this matters

Every material an engineer specifies, every reagent a lab technician measures and every blood test a medical laboratory scientist runs comes down to atoms. The periodic table is the reference sheet that packs the most useful facts about every element into one page. If you can read it fluently, the rest of this course (bonding, reactions, moles, pH) becomes arithmetic instead of memorisation.

## What an atom is made of

An atom has a tiny, dense **nucleus** surrounded by **electrons**.

| particle | charge | where it is | mass (approx.) |
|---|---|---|---|
| proton | +1 | nucleus | 1 u |
| neutron | 0 | nucleus | 1 u |
| electron | −1 | around the nucleus | about 1/1836 u |

`u` is the atomic mass unit (also called a dalton, Da). Electrons are so light that an atom's mass is almost entirely its protons and neutrons.

Two numbers describe an atom:

- **Atomic number (Z)**: the number of protons. It defines the element. Every carbon atom has 6 protons; an atom with 7 protons is nitrogen, never carbon.
- **Mass number (A)**: protons + neutrons. So neutrons = A − Z.

A neutral atom has as many electrons as protons, so its charges cancel.

## Isotopes and atomic mass

Atoms of the same element can have different numbers of neutrons. These are **isotopes**. Carbon-12 has 6 protons and 6 neutrons; carbon-14 has 6 protons and 8 neutrons. Same element, same chemistry, different mass.

The **atomic mass** printed on the table (for example 35.45 for chlorine) is a weighted average over the isotopes found in nature. Chlorine is about 75.8% chlorine-35 (mass 34.969 u) and 24.2% chlorine-37 (mass 36.966 u):

```text
0.7576 × 34.969 + 0.2424 × 36.966 ≈ 26.49 + 8.96 = 35.45 u
```

That is why atomic masses are rarely whole numbers, and why you never see a "35.45" mass number on a single atom.

## Reading an element tile

A typical tile for carbon shows:

```text
   6        <- atomic number Z
   C        <- symbol
 Carbon     <- name
 12.011     <- average atomic mass (u, or g/mol)
```

Worked example: an atom of **sodium-23**.

- The table gives Na, Z = 11, so 11 protons.
- Mass number 23, so 23 − 11 = **12 neutrons**.
- Neutral, so **11 electrons**.

Worked example: the ion **Cl⁻** (chloride) from chlorine-35.

- Z = 17, so 17 protons; 35 − 17 = 18 neutrons.
- The 1− charge means one extra electron: 17 + 1 = **18 electrons**.

A positive ion has lost electrons: Mg²⁺ has 12 protons and 10 electrons.

## How the table is organised

- **Periods** are the rows (1–7). Moving left to right, Z goes up by one each step.
- **Groups** are the columns (1–18). Elements in a group have the same number of outer (**valence**) electrons for the main-group elements, so they behave alike.
- **Metals** sit on the left and in the middle, **nonmetals** in the upper right, and **metalloids** (such as silicon) along the staircase between them.

Groups worth knowing by name:

| group | name | typical behaviour | common ion |
|---|---|---|---|
| 1 | alkali metals (Li, Na, K) | very reactive metals | +1 |
| 2 | alkaline earth metals (Mg, Ca) | reactive metals | +2 |
| 17 | halogens (F, Cl, Br, I) | reactive nonmetals | −1 |
| 18 | noble gases (He, Ne, Ar) | very unreactive | none |

Main-group elements gain or lose electrons to reach the same electron count as the nearest noble gas. That is why sodium forms Na⁺ (it loses 1 to match neon) and oxygen forms O²⁻ (it gains 2 to match neon).

### Trends you can predict

- **Atomic radius** grows down a group (more electron shells) and shrinks across a period (more protons pull the electrons in).
- **Electronegativity**, how strongly an atom pulls on shared electrons, rises toward the upper right. Fluorine is the most electronegative element. You will use this in the next lesson to predict bond types.

## Common mistakes

- **Confusing atomic number and mass.** The small whole number is Z; the decimal is the average mass.
- **Changing protons to make an ion.** Ions gain or lose *electrons*. Change the protons and you have a different element.
- **Rounding the atomic mass and calling it the mass number of every atom.** Chlorine atoms have mass number 35 or 37; none has 35.45.
- **Expecting noble gases to form ions.** Under ordinary conditions they almost never react.

## Where this goes next

Chemistry is a gateway to lab, engineering and materials careers, but the platform teaches the knowledge, not the hands-on lab skill. Chemist roles typically require a bachelor's degree in chemistry (ACS-approved programmes are often preferred) with real laboratory coursework; medical laboratory scientists also need an accredited programme and certification. See the [American Chemical Society](https://www.acs.org/careers.html) and [ASCP Board of Certification](https://www.ascp.org/content/board-of-certification). Checked on 2026-10-23.

## Quick recap

- Protons define the element (Z); protons + neutrons = mass number (A).
- Isotopes differ in neutrons; the table lists a weighted average mass.
- Ions form by gaining or losing electrons, never protons.
- Groups share valence electrons and behaviour; trends run across periods and down groups.

## Try it

Open the **Do** tab and answer the quiz: you will read element data off the table and count particles in atoms and ions.


Read an element's atomic number, symbol and atomic mass from the periodic table and work out its protons, neutrons and electrons.

Atoms, elements, and the periodic table

The Periodic Table: Crash Course Chemistry #4

The periodic table | Atoms, elements, and the periodic table | High school chemistry | Khan Academy

## Why this matters

Table salt melts at 801 °C and dissolves in water; candle wax melts in your hand and does not. Steel conducts electricity; plastic insulates it. These differences come from how atoms are held together. Predicting the bond type from two element symbols lets you guess whether a substance will dissolve, conduct, melt easily or stay solid, which is everyday reasoning in labs, materials work and engineering.

## The core idea: who holds the electrons?

A bond forms because atoms are more stable when they reach a full outer shell of electrons. They get there by **transferring** electrons or **sharing** them. Which one happens depends on how unequally the two atoms pull on electrons, measured by **electronegativity** (Pauling scale, from about 0.7 to 3.98).

| bond type | what happens to the electrons | typical partners |
|---|---|---|
| ionic | transferred from one atom to the other, forming ions | metal + nonmetal |
| polar covalent | shared unequally; one end slightly negative (δ−) | two different nonmetals |
| nonpolar covalent | shared (almost) equally | identical nonmetals, or C–H |
| metallic | outer electrons spread over the whole metal | metal + metal |

### A rule of thumb with electronegativity difference (ΔEN)

```text
ΔEN < 0.4         nonpolar covalent
0.4 ≤ ΔEN < ~1.8  polar covalent
ΔEN ≥ ~1.8        mostly ionic
```

The cut-offs are approximate: bonding is a spectrum, and textbooks place the ionic boundary anywhere from about 1.7 to 2.0. The metal + nonmetal shortcut usually gives the same answer.

## Worked examples

Electronegativity values: Na 0.93, Mg 1.31, H 2.20, C 2.55, N 3.04, O 3.44, Cl 3.16, F 3.98.

**NaCl (sodium chloride).** ΔEN = 3.16 − 0.93 = 2.23, and it is a metal with a nonmetal: **ionic**. Sodium gives its single valence electron to chlorine, making Na⁺ and Cl⁻. The ions stack into a crystal lattice held by strong attractions in every direction, which is why salt is hard, brittle and has a high melting point. Molten or dissolved, the ions can move, so it conducts electricity.

**MgO (magnesium oxide).** ΔEN = 3.44 − 1.31 = 2.13: **ionic**. Mg²⁺ and O²⁻ carry double charges, so the attraction is even stronger; MgO melts at about 2,850 °C and is used to line furnaces.

**H₂O (water).** O–H: ΔEN = 3.44 − 2.20 = 1.24: **polar covalent**. Oxygen pulls the shared electrons toward itself, so oxygen is δ− and each hydrogen δ+. Because the molecule is bent, those partial charges do not cancel, and water is a polar molecule. That polarity is why water dissolves salt and sugar so well.

**CH₄ (methane).** C–H: ΔEN = 2.55 − 2.20 = 0.35: **nonpolar covalent**. Methane, like wax and oils, is mostly C–H bonds, which is why it does not mix with water.

**Cl₂ (chlorine gas).** Two identical atoms, ΔEN = 0: **nonpolar covalent**.

**CO₂ (carbon dioxide).** Each C=O bond is polar (ΔEN = 0.89), but the molecule is straight (O=C=O), so the two pulls cancel. CO₂ has polar bonds but is a **nonpolar molecule**. Shape matters.

## Single, double and triple bonds

Atoms can share one, two or three pairs of electrons:

- single: H–H, C–H
- double: O=O, C=O
- triple: N≡N (nitrogen gas, very hard to break, which is why N₂ is unreactive)

More shared pairs means a shorter, stronger bond.

## Properties follow from bond type

| | ionic compounds | covalent molecules | metals |
|---|---|---|---|
| melting point | high | usually low | usually high |
| conducts electricity | when molten or dissolved | rarely | yes, as solids |
| in water | many dissolve | polar ones dissolve; nonpolar ones do not | do not dissolve |

"Like dissolves like" is the practical summary: polar and ionic substances mix with water; nonpolar substances mix with oils and other nonpolar solvents.

## Common mistakes

- **Saying ionic compounds are made of molecules.** Solid NaCl is a lattice of ions; "NaCl" is a ratio, not a molecule.
- **Treating polar bonds as a polar molecule.** CO₂ and CCl₄ have polar bonds but symmetric shapes, so they are nonpolar.
- **Using the ΔEN cut-off as a hard law.** It is a guide. Hydrogen fluoride (ΔEN 1.78) is covalent even though it sits near the boundary.
- **Forgetting metallic bonding.** Two metals do not form an ionic bond; they form an alloy held by metallic bonding.

## Quick recap

- Bonds form by transferring (ionic) or sharing (covalent) electrons.
- Electronegativity difference predicts the type; metal + nonmetal is usually ionic.
- Polar bonds in a symmetric molecule can cancel out.
- Bond type explains melting point, conductivity and solubility.

## Try it

Open the **Do** tab and take the quiz: you will classify bonds from element pairs and predict properties from bond type.


Use electronegativity and position on the periodic table to predict whether a bond is ionic, polar covalent or nonpolar covalent.

Chemical bonding

Atomic Hook-Ups - Types of Chemical Bonds: Crash Course Chemistry #22

The Chemical Bond: Covalent vs. Ionic and Polar vs. Nonpolar

## Why this matters

Gases are everywhere in technical work: compressed air in a plant, carbon dioxide in a cell-culture incubator, nitrogen purges in a chemistry lab, natural gas in a building. Engineers size tanks and pipes with gas calculations, and lab staff convert gas volumes to amounts every day. The ideal gas law is the one equation that ties pressure, volume, temperature and amount together.

## The three everyday states

Matter is made of particles that are always moving. The state depends on how strongly they attract each other compared with how much energy they have.

| state | particle arrangement | shape and volume |
|---|---|---|
| solid | packed tightly, vibrating in place | fixed shape, fixed volume |
| liquid | close together but able to slide past each other | takes the container's shape, fixed volume |
| gas | far apart, moving fast and randomly | fills the whole container |

Adding heat gives particles more energy, so substances **melt** (solid to liquid) and **boil** or **evaporate** (liquid to gas). Removing heat reverses this: **condensation** and **freezing**. Some solids go straight to gas (**sublimation**), like dry ice (solid CO₂). During a change of state the temperature stays constant while the energy goes into breaking attractions: an ice-water mix stays at 0 °C until the ice is gone.

## Temperature must be in kelvin

Gas laws only work with an absolute temperature scale, where zero means no thermal motion:

```text
T(K) = T(°C) + 273.15
25 °C  -> 298.15 K
0 °C   -> 273.15 K
```

If you plug in Celsius, 0 °C would make the volume of every gas zero, which is clearly wrong.

## The ideal gas law

```text
P V = n R T
```

- `P` pressure, `V` volume, `n` amount in moles, `T` temperature in kelvin
- `R` the gas constant. Its value depends on the units you use:

| units of P and V | R |
|---|---|
| kPa and litres | 8.314 L·kPa/(mol·K) |
| atm and litres | 0.08206 L·atm/(mol·K) |
| Pa and m³ | 8.314 J/(mol·K) |

Pick one row and keep every quantity in those units. Unit mismatches are the most common source of wrong answers.

### Worked example 1: amount of gas in a balloon

A balloon holds 5.0 L of helium at 100 kPa and 25 °C. How many moles?

```text
T = 25 + 273.15 = 298.15 K
n = PV / (RT) = (100 × 5.0) / (8.314 × 298.15)
  = 500 / 2478.8 ≈ 0.202 mol
```

### Worked example 2: pressure in a sealed container

2.0 mol of nitrogen is sealed in a 10.0 L rigid container at 25 °C.

```text
P = nRT / V = (2.0 × 8.314 × 298.15) / 10.0 ≈ 496 kPa
```

That is almost five times normal atmospheric pressure (101.325 kPa), which is why gas cylinders are thick steel and are handled only after training.

### Worked example 3: the molar volume

At 0 °C and 1 atm (101.325 kPa), one mole of any ideal gas occupies

```text
V = nRT / P = (1 × 8.314 × 273.15) / 101.325 ≈ 22.4 L
```

You may also see 22.7 L: that is for the IUPAC standard of 0 °C and 100 kPa. Always check which "standard conditions" a source means.

## The simpler gas laws are special cases

Hold two of the four quantities fixed and PV = nRT gives the classic laws:

- **Boyle:** at constant n and T, `P₁V₁ = P₂V₂` (squeeze a gas and its pressure rises).
- **Charles:** at constant n and P, `V₁/T₁ = V₂/T₂` (warm a balloon and it expands).
- **Gay-Lussac:** at constant n and V, `P₁/T₁ = P₂/T₂` (a heated sealed can builds pressure).

## When "ideal" stops being true

The ideal gas law assumes the particles take up no space and do not attract each other. That is a good approximation for most gases near room temperature and atmospheric pressure. It breaks down at **high pressure** and **low temperature**, near the point where the gas would condense. Engineers then use real-gas equations (such as van der Waals) or tables.

## Safety note

Compressed-gas cylinders, cryogenic liquids and pressure vessels are real hazards. This lesson is awareness level only: handling them requires your employer's or institution's training and supervision. OSHA's [compressed gases page](https://www.osha.gov/compressed-gas-equipment) is a starting point. Checked on 2026-10-23.

## Common mistakes

- **Celsius in the equation.** Always convert to kelvin first.
- **Mixing units.** kPa with the atm value of R gives an answer off by about 100 times.
- **Forgetting mL to L.** 250 mL is 0.250 L.
- **Applying it to liquids or solids.** PV = nRT is for gases only.

## Quick recap

- Solids, liquids and gases differ in particle spacing and motion.
- PV = nRT links pressure, volume, amount and absolute temperature.
- Match R to your units; one mole of gas is about 22.4 L at 0 °C and 1 atm.

## Try it

Open the **Do** tab: you will finish a Python function that calculates moles of gas from pressure, volume and temperature, then answer the quiz.


Explain solids, liquids and gases in terms of particle motion and use the ideal gas law PV = nRT with temperatures in kelvin.

States of matter and gases

The Ideal Gas Law: Crash Course Chemistry #12

Ideal Gas Problems: Crash Course Chemistry #13

Atomic structure and the periodic table, ionic and covalent bonding, and the states of matter with the ideal gas law.