Matter –
anything that occupies space and have mass
Three states of matter:
solid – has definite shape and definite volume. In solid state atoms pack closely together in
fixed location. Neighboring atoms vibrate or oscillate, but do not move around each other.
Solid mater may be crystalline (atoms and molecules are arranged in long-range, repeating order) or
amorphous (atoms and molecules are not arranged in long-range, repeating order).
liquid – has definite volume and indefinite shape. Liquid assume the shape of their container
because atoms are free to move around each other.
gas – has indefinite shape and indefinite volume. Atoms are separated by large (in atomic
scale!) distances. Because atoms or molecules of a gas are not in contact with one another, gases are
compressible.
Energy –
capacity to do work. There are two states (=work capacities) of energy: potential and kinetic energy.
Kinetic energy – energy in action. Examples: constant movement of atoms; bouncing ball.
Potential energy – stored, or inactive, energy. When potential energy is released, it converts to
kinetic energy. Example: unused battery; ATP molecule.
Forms of energy:
chemical energy – energy stored in chemical bonds; it is liberated when chemical bonds break
in chemical reaction.
electrical energy – results from movement of charged particles.
mechanical energy – energy directly involved in movement.
electromagnetic (=radiant) energy – travels in waves (light=photons; x-ray.)
Composition of matter
Elements – unique substances that cannot be broken by chemical means. Each element is composed of
a small particle – an atom – that gives them their unique properties.
Physical properties of an element can with our senses or my measuring them (the boiling of an
element is an example of measuring its physical properties).
Chemical properties can be detected through chemical reactions. Chemical properties determine
how atoms of an element interact with each other.
Atoms – smallest, “indivisible” particles of matter.
Atomic structure
An atom consist of a nucleus with its protons and neutrons, and orbiting around the nucleus electron.
The number of protons in nucleus is always equal to the number of electrons.
Proton – a positively charged particle of atom's nucleus.
Neutron – a particle inside the nucleus with no charge.
A mass of an atom is roughly equal to a combined mass of its protons and neutrons.
Electron – a negatively charged particle of an atom. An electron is so small and light compared
to a proton that its mass is disregarded in measuring atomic mass.
Positive charge of protons and negative charge of electrons cancel each other out making atom
electrically neutral.
An element properties are determined by the number of protons in the nucleus of its atom. For
example, hydrogen atom has one proton and helium atom has two protons. Elements hydrogen and
helium will have different physical and chemical properties. The number of neutrons in nucleus,
however, may vary.
Isotops –
atoms of an element that have the same number of protons, but different number of neutrons in the
nucleus.
Molecules –
combination of two or more atoms.
If two or more atoms of the same element are combined, they form a molecule of that element.
Examples: O2 – oxygen gas, N2 – nitrogen gas.
Compounds –
combination of two or more atoms of different elements. Examples: H2O – water, CO – carbon
monoxide (gas).
Mixtures –
substances composed of two or more elements that are physically intermixed. Example: air.
Solutions –
homogenous mixtures of components. Example: seawater. The substance that is present in greater
amount in a solution is called solvent. The substance that is present in smaller amount in a solution is
called solute. Example: salt in water → water – solvent; salt – solute.
Colloids –
heterogenous mixtures, translucent or milky, in which particles do not settle out in a solution. Example:
milk.
Some colloids are able to change from fluid state (sol) to more solid state (gel); this is called sol-gel
transformation. Example: Jell-O.
Suspensions –
heterogenous mixtures in which particles do not settle out in a solvent. Example: blood.
Distinguishing mixtures from compounds
In compounds, chemical properties of separate elements change due to chemical bonding between
atoms. In mixtures, components do not create chemical bonds between each other, therefore, their
chemical properties remain the same.
Chemical bonds
When atoms or molecules combine, they are held together by chemical bonds. Chemical bond is
electrical force that links electrons of separate atoms.
Ionic bonds –
transfer of one or more electrons from one atom to another. The atom that loses its electron is called
cation (=positive ion) because it becomes positively charged. The atom that gains an electron is called
anion (=anion) because its net charge becomes negative. Example: NaCl → Na+ClCovalent bonds –
atoms in a compound do not lose or gain electrons. Instead their electrons are shared between two
atoms.
Non-polar covalent bonds – electrons are equally shared between atoms of a molecule. The molecule
created by non-polar covalent charge will be electrically balanced. Example: O2
Polar covalent bonds – electrons are not shared equally between the atoms. The molecule created by
polar bonds will have positive and negative poles due to different electron-attracting abilities of its
atoms. Example H2O.
Hydrogen bonds –
intramolecular bonds. Hydrogen bonds are attractions between the hydrogen atom of one molecule and
the electronegative pole of another molecule (usually oxygen or nitrogen). Water molecules form
hydrogen bonds, and these bonds give water its fluid properties. Hydrogen bonds are also responsible
for surface tension – tendency of water molecules to cling together and form a film.
Chemical reactions –
occur whenever chemical bonds are rearranged, formed, or broken.
Synthesis (=combination) reaction – atoms or molecules combine together to form a new,
larger and more complex, molecule. Example: N + N → N2
Decomposition reaction – large, complex molecules are broken down into smaller molecules or
atoms. Example: 2H2O → 2H + O2
Replacement (=exchange) reaction – molecules or atoms are rearranged into new molecules. In
replacement reactions, both synthesis and decomposition will occur.
Example: Mg + 2H2O → Mg(OH)2 + H2
Factors influencing chemical reaction
Temperature –
when temperature increases, the kinetic energy of substance inreases. The particles will collide with
more force at higher pace. This will speed up chemical reaction.
Concentration –
the higher the number of particles in a solution, the higher is the chance of successful collisions
between them; therefore, chemical reaction will progress quicker.
Particle size –
smaller particles move faster and collide more frequently than large particle.
Catalyst –
a substance that speeds up the chemical reaction without being chemically rearranged itself. Example:
enzymes are biological catalysts.