Vocabulary for Standard 4 Kinetics




3rd Law of Thermodynamics    
absolute temperature    
absolute zero    

Amorphous solid

amorphous solid.



A solid that does not have a repeating, regular three-dimensional arrangement of atoms, molecules, or ions.

Average Kinetic Energy    



Boiling, a type of phase transition, is the rapid vaporization of a liquid, which typically occurs when a liquid is heated to its boiling point, the temperature at which the vapor pressure of the liquid is equal to the pressure exerted on the liquid by the surrounding atmospheric pressure. Thus, a liquid may also boil when the pressure of the surrounding atmosphere is sufficiently reduced, such as the use of a vacuum pump or at high altitudes. Boiling occurs in three characteristic stages, which are nucleate, transition and film boiling. These stages generally take place from low to high surface temperatures, respectively.

Boiling point

(bp) standard boiling point; normal boiling point.



The temperature at which the vapor pressure of a liquid is equal to the external pressure on the liquid. The standard boiling point is the temperature at which the vapor pressure of a liquid equals standard pressure.



An insulated vessel for measuring the amount of heat absorbed or released by a chemical or physical change.


(C) Celsius temperature scale; Celsius scale.



A common but non-SI unit of temperature, defined by assigning temperatures of 0C and 100C to the freezing and boiling points of water, respectively.









1. The conversion of a gas into a liquid is called condensation. Condensation usually occurs when a gas is cooled below its boiling point. 2. A reaction that involves linking of two molecules with the elimination of water (or another small molecule).




Critical point

critical state.



State at which two phases of a substance first become indistinguishable. For example, at pressures higher than 217.6 atm andtemperatures above 374C, the meniscus between steam and liquid water will vanish; the two phases become indistinguishable and are referred to as a supercritical fluid.

Critical pressure




The pressure at the critical point.

Critical temperature




The temperature at the critical point. A gas above the critical temperature will never condense into a liquid, no matter how much pressure is applied. Most substances have a critical temperature that is about 1.5 to 1.7 times the standard boiling point, in kelvin.


Crystal structure



Crystalline solids




A solid that has a repeating, regular three-dimensional arrangement of atoms, molecules, or ions.








The mixing of two substances caused by random molecular motions. Gases diffuse very quickly; liquids diffuse much more slowly, and solids diffuse at very slow (but often measurable) rates. Molecular collisions make diffusion slower in liquids and solids.





Gas molecules in a container escape from tiny pinholes into a vacuum with the same average velocity they have inside the container. They also move in straight-line trajectories through the pinhole.

Elastic collision




endothermic reaction; endothermic process. .


A process that absorbs heat. The enthalpy change for an endothermic process has a positive sign.


Enthalpy change

(H) enthalpy change.



Enthalpy (H) is defined so that changes in enthalpy ( DeltaH) are equal to the heat absorbed or released by a process running at constant pressure. While changes in enthalpy can be measured using calorimetry, absolute values of enthalpy usually cannot be determined. Enthalpy is formally defined as H = U + PV, where U is the internal energy, P is the pressure, and V is the volume.

Enthalpy of combustion

( DeltaHc) heat of combustion.


The change in enthalpy when one mole of compound is completely combusted. All carbon in the compound is converted to CO2(g), all hydrogen to H2O( ell), all sulfur to SO2(g), and all nitrogen to N2(g).


Enthalpy of reaction

( DeltaHrxn) heat of reaction.



The heat absorbed or released by a chemical reaction running at constant pressure.

Equilibrium vapor pressure





To convert a liquid into a gas.





Conversion of a liquid into a gas.


exothermic reaction; exothermic process.


A process that releases heat. The enthalpy change for an exothermic process is negative. Examples of exothermic processes are combustion reactions and neutralization reactions











Freezing point

(mp) standard melting point; normal melting point; melting point.

The temperature at which the vapor pressure of a liquid is equal to the vapor pressure of the corresponding solid form. The liquid and solid forms can coexist at equilibrium at the freezing point. The standard melting point is the melting point at standard pressure.

Graph: T vs phase of H2O






Heat is a transfer of energy that occurs when objects with different temperatures are placed into contact. Heat is a process, not a property of a material.

Hesss law

law of constant heat summation; Hess's law of heat summation.

The heat released or absorbed by a process is the same no matter how many steps the process takes. For example, given a reaction A rightarrowB, Hess's law says that DeltaH for the reaction is the same whether the reaction is written as A rightarrowC rightarrowB or as A rightarrowB. This is the same as writing that DeltaH(A rightarrowB) = DeltaH(A rightarrowC) + DeltaH(C rightarrowB).

internal energy    



The SI unit of energy, equal to the work required to move a 1 kg mass against an opposing force of 1 newton. 1 J = 1 kg m2 s-2 = 4.184 calories.



The SI base unit of temperature, defined by assigning 273.16 K to the temperature at which steam, ice, and water are at equilibrium (called the triple point of water). The freezing point of water is 273.15 K.

Kinetic energy


The energy an object possesses by virtue of its motion. An object of mass m moving at velocity v has a kinetic energy of mv2.

Kinetic molecular model    

Kinetic molecular theory






Melting point



Normal boiling point

See Boiling point



in phase; out of phase; wave phase.

1.A phase is a part of a sample of matter that is in contact with other parts but is separate from them. Properties within a phase are homogeneous (uniform). For example, oil and vinegar salad dressing contains two phases: an oil-rich liquid, and a vinegar-rich liquid. Shaking the bottle breaks the phases up into tiny droplets, but there are still two distinct phases. 2. In wave motion, phase is the fraction of a complete cycle that has passed a fixed point since the current cycle began. The phase is often expressed as an angle, since a full cycle is 360^deg; (2 pi). Two waves are "in phase" if the peaks of one wave align with the peaks of the other; they are "out of phase" if the peaks of one wave align with the troughs of the other.

Phase diagram

phase map.

A map that shows which phases of a sample are most stable for a given set of conditions. Phases are depicted as regions on the map; the borderlines between regions correspond to conditions where the phases can coexist in equilibrium.

potential energy    
standard temperature and pressure   One atmosphere and 273 K.


sublimate; sublimating.

Conversion of a solid directly into a gas, without first melting into a liquid.



Temperature is an intensive property associated with the hotness or coldness of an object. It determines the direction of spontaneous heat flow (always from hot to cold).

Thermochemical equation


An compact equation representing a chemical reaction that describes both the stoichiometry and the energetics of the reaction. For example, the thermochemical equation CH4(g) + 2 O2(g) rightarrowCO2(g) + 2 H2O(g), DeltaH = -2220 kJ means "When one mole of gaseous CH4 is burned in two moles of oxygen gas, one mole of CO2 gas and 2 moles of steam are produced, and 2220 kilojoules of heat are released."



The study of heat absorbed or released during chemical changes.

Transition states


In a chemical reaction, the reagents have to join together into a great big blob before they can fall back apart into the products.  This great big blob is called the activated complex

Triple point


The temperature and pressure at which the solid, liquid, and gaseous forms of a substance are at equilibrium.

Vapor Pressure

vapour pressure.

The partial pressure of a gas in equilibrium with a condensed form (solid or liquid) of the same substance.


See evaporate


Volatile liquids


A solid or liquid material that easily vaporizes. A material with a significant vapor pressure.