Basic Chemistry: Matter, Atoms, Molecules, and Mixtures
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74 Terms
Matter
Anything that has mass and takes up space.
Atoms
The smallest units of matter that make up all substances.
Molecule
Two or more atoms chemically bonded together.
Element
A pure substance made of only one type of atom.
Compound
A pure substance made of two or more different elements chemically combined in a fixed ratio.
Pure Substance
A substance made of only one kind of particle (element or compound) with definite properties.
Examples of Pure Substances
Gold (Au), oxygen gas (O₂), water (H₂O), and salt (NaCl).
Mixture
A physical combination of two or more substances that are not chemically bonded and can be separated.
Homogeneous Mixture
A mixture that looks uniform throughout, like salt water or air.
Heterogeneous Mixture
A mixture that is not uniform and where you can see different parts, like salad or cereal with milk.
Examples of Homogeneous Mixtures
Salt water, iced tea, air.
Examples of Heterogeneous Mixtures
Salad, trail mix, sand and water.
Diatomic Molecule
Two atoms of the same element bonded together, like H₂, O₂, N₂, or Cl₂.
Allotrope
Different forms of the same element with atoms arranged differently, like O₂ and O₃ or graphite and diamond.
Can Compounds Be Separated Physically?
No, only chemically.
Can Mixtures Be Separated Chemically?
No, only physically.
Is Water an Element, Compound, or Mixture?
Compound (H₂O).
Is Salt an Element, Compound, or Mixture?
Compound (NaCl).
Is Air an Element, Compound, or Mixture?
Mixture.
Is Gold an Element, Compound, or Mixture?
Element.
Is Lemonade an Element, Compound, or Mixture?
Mixture (homogeneous).
Is Sand an Element, Compound, or Mixture?
Mixture (heterogeneous).
Physical change
A change that affects appearance or state but not the chemical composition.
Examples of physical changes
Melting, freezing, cutting, dissolving.
Chemical change
A change that forms a new substance with new properties.
Examples of chemical changes
Rusting, burning, reacting, cooking food.
Law of Conservation of Matter
Matter cannot be created or destroyed, only rearranged.
Energy
The ability to do work or cause change.
Types of energy
Kinetic energy and potential energy.
Kinetic energy
Energy of motion; particles moving faster have more kinetic energy.
Potential energy
Stored energy based on position or chemical bonds.
Thermal energy
The total kinetic energy of particles in a substance.
Heat
The transfer of thermal energy from a hotter object to a colder one.
Temperature
A measure of the average kinetic energy of particles in a substance.
Difference between temperature and heat
Temperature measures how hot something is; heat is energy transfer caused by temperature difference.
Exothermic reaction
A reaction that releases heat and feels hot; ΔH is negative.
Endothermic reaction
A reaction that absorbs heat and feels cold; ΔH is positive.
Activation energy
The energy needed to start a chemical reaction.
Activated complex
The temporary "in-between" substance formed during a reaction.
Heat of reaction (ΔH)
The difference in energy between products and reactants.
q = mcΔT
Calculates the amount of heat energy absorbed or released when a substance changes temperature.
q in q = mcΔT
Heat energy, measured in joules (J).
m in q = mcΔT
Mass, measured in grams (g).
c in q = mcΔT
Specific heat capacity, measured in J/g°C.
ΔT in q = mcΔT
Change in temperature (final temperature - initial temperature).
When to use q = mcΔT
When the substance is heating or cooling but not changing phase.
q = mHf
Used when the substance is melting or freezing.
q = mHv
Used when the substance is boiling or condensing.
Hf (heat of fusion)
The energy required to melt or freeze 1 gram of a substance.
Hv (heat of vaporization)
The energy required to boil or condense 1 gram of a substance.
Heat of fusion for water
334 J/g.
Heat of vaporization for water
2260 J/g.
Specific heat of water
4.18 J/g°C.
Positive q value
Heat is absorbed (endothermic).
Negative q value
Heat is released (exothermic).
Energy changes during temperature changes
Kinetic energy.
Energy changes during phase changes
Potential energy.
Temperature during a phase change
No, it stays constant while potential energy changes.
Energy during melting
Absorbed (endothermic).
Energy during freezing
Released (exothermic).
Energy during boiling
Absorbed (endothermic).
Energy during condensation
Released (exothermic).
Sloped lines in a heating curve
Represent temperature changes, so kinetic energy changes.
Flat lines in a heating curve
Represent phase changes, so potential energy changes.
Equation for temperature changes
q = mcΔT.
Units for heat (q)
Joules (J).
Units for specific heat (c)
J/g°C.
Units for Hf and Hv
J/g.
Specific heat capacity
Measures how much energy it takes to raise 1 gram of a substance by 1°C.
Substance with high specific heat
Heats up and cools down slowly.
Substance with low specific heat
Heats up and cools down quickly.
Celsius to Kelvin conversion
K = °C + 273.
Celsius to Fahrenheit conversion
°F = (°C × 9/5) + 32.
Chemical change indicators
Temperature change, bubbles of gas, change in smell or taste, change in melting or boiling point