Heat and Temperature

Difference Between Heat and Temperature

Difference Between Heat and Temperature

– Heat is one form of energy.

– Many forms of energy can be interconverted and that in chemical processes, chemical energy is converted to heat energy or vice versa.

– The amount of heat a process uses (endothermic) or gives off (exothermic) can tell us a great deal about that process.

– For this reason it is important for us to be able to measure the intensity of heat.

– Temperature measures the intensity of heat, the “hotness” or “coldness” of a body.

– A piece of metal at 100°C feels hot to the touch, whereas an ice cube at 0°C feels cold. Why?

– Because the temperature of the metal is higher, and that of the ice cube lower, than body temperature.

– Heat is a form of energy that always flows spontaneously from a hotter body to a colder body—never in the reverse direction.

– Temperatures can be measured with mercury-in-glass thermometers.

– A mercury thermometer consists of a reservoir of mercury at the base of a glass tube, open to a very thin (capillary) column extending upward.

Mercury expands more than most other liquids as its temperature rises.

– As it expands, its movement up into the evacuated column can be seen.

Celsius temperature scale

– Anders Celsius (1701–1744), a Swedish astronomer, developed the Celsius temperature scale, formerly called the centigrade temperature scale.

– When we place a Celsius thermometer in a beaker of crushed ice and water, the mercury level stands at exactly 0°C, the lower reference point.

– In a beaker of water boiling at one atmosphere pressure, the mercury level stands at exactly 100°C, the higher reference point.

– There are 100 equal steps between these two mercury levels.

– They correspond to an interval of 100 degrees between the melting point of ice and the boiling point of water at one atmosphere.

– Figure (1) shows how temperature marks between the reference points are established.

Fahrenheit temperature scale

– In the United States, temperatures are frequently measured on the temperature scale devised by Gabriel Fahrenheit (1686–1736), a German instrument maker.

– On this scale the freezing and boiling points of water are defined as 32°F and 212°F, respectively.

Kelvin temperature scale

– In scientific work, temperatures are often expressed on the Kelvin (absolute) temperature scale.

the zero point of the Kelvin temperature scale is derived from the observed behavior of all matter.

Relationships among Celsius, Fahrenheit, Kelvin temperature scales

– Relationships among the three temperature scales are illustrated in Figure (2).

Heat and Temperature

– Between the freezing point of water and the boiling point of water, there are 100 steps (°C or kelvins, respectively) on the Celsius and Kelvin scales.

– Thus the “degree” is the same size on the Celsius and Kelvin scales.

– But every Kelvin temperature is 273.15 units above the corresponding Celsius temperature.

– The relationship between these two scales is as follows:

– In the SI system, “degrees Kelvin” are abbreviated simply as K rather than °K and are called kelvins.

– Any temperature change has the same numerical value whether expressed on the Celsius scale or on the Kelvin scale.

– For example, a change from 25°C to 59°C represents a change of 34 Celsius degrees.

– Converting these to the Kelvin scale, the same change is expressed as (273 + 25) = 298 K to (59 + 273) = 332 K, or a change of 34 kelvins.

– Comparing the Fahrenheit and Celsius scales, we find that the intervals between the same reference points are 180 Fahrenheit degrees and 100 Celsius degrees, respectively.

– Thus a Fahrenheit degree must be smaller than a Celsius degree. It takes 180 Fahrenheit degrees to cover the same temperature interval as 100 Celsius degrees.

– From this information, we can construct the unit factors for temperature changes:

Heat and Temperature

Heat and Temperature

– These are often remembered in abbreviated form:

Heat and Temperature

– But the starting points of the two scales are different, so we cannot convert a temperature on one scale to a temperature on the other just by multiplying by the unit factor.

– In converting from °F to °C, we must subtract 32 Fahrenheit degrees to reach the zero point on the Celsius scale (Figure 2).

Solved problems on Heat and Temperature

Example (1): When the temperature reaches “100.°F in the shade,” it’s hot. What is this temperature on the Celsius scale?


– We use the relationship:

to carry out the desired conversion.


Example (2): When the absolute temperature is 400 K, what is the Fahrenheit temperature?


We first use the relationship:

__ ? °C= K + 273°

to convert from kelvins to degrees Celsius, then we carry out the further conversion from degrees Celsius to degrees Fahrenheit.


Heat and Temperature

Heat Transfer and The Measurements of Heat

– Chemical reactions and physical changes occur with either the simultaneous evolution of heat (exothermic processes) or the absorption of heat (endothermic processes).

– The amount of heat transferred in a process is usually expressed in joules or in calories.

– The SI unit of energy and work is the joule (J), which is defined as 1 kg.m2/s2.

– The kinetic energy (KE) of a body of mass m moving at speed v is given by 1/2 mv2.

– A 2-kg object moving at one meter per second has KE = 1/2(2 kg)(1 m/s)2 =1 kg.m2/s2 = 1 J.

– You may find it more convenient to think in terms of the amount of heat required to raise the temperature of one gram of water from 14.5°C to 15.5°C, which is 4.184 J.

– One calorie is defined as exactly 4.184 J.

– The so-called “large calorie,” used to indicate the energy content of foods, is really one kilocalorie, that is, 1000 calories. We shall do most calculations in joules.

The specific heat

The specific heat of a substance is the amount of heat required to raise the temperature of one gram of the substance one degree Celsius (also one kelvin) with no change in phase.

– Changes in phase (physical state) absorb or liberate relatively large amounts of energy .

– The specific heat of each substance, a physical property, is different for the solid, liquid, and gaseous phases of the substance.

– For example, the specific heat of ice is 2.09 J/g.°C near 0°C; for liquid water it is 4.18 J/g.°C; and for steam it is 2.03 J/g.°C near 100°C.

– The specific heat for water is quite high.

Heat and Temperature

The heat capacity

– The heat capacity of a body is the amount of heat required to raise its temperature 1°C.

– So, The heat capacity of a body is its mass in grams times its specific heat.

– The heat capacity refers to the mass of that particular body, so its units do not include mass. The units are J/°C or J.°C-1.

Solved problems on The specific heat

Example (1): How much heat, in joules, is required to raise the temperature of 205 g of water from 21.2°C to 91.4°C?


– The specific heat of a substance is the amount of heat required to raise the temperature of 1 g of substance 1°C:

– We can rearrange the equation so that:

(Amount of heat) = (mass of substance) (specific heat) (temperature change)

– Alternatively, we can use the unit factor approach.


By the unit factor approach,

– All units except joules cancel. To cool 205 g of water from 91.4°C to 21.2°C, it would be necessary to remove exactly the same amount of heat, 60.2 kJ

– When two objects at different temperatures are brought into contact, heat flows from the hotter to the colder body (Figure 3); this continues until the two are at the same temperature.

Heat and Temperature

– We say that the two objects are then in thermal equilibrium.

– The temperature change that occurs for each object depends on the initial temperatures and the relative masses and specific heats of the two materials.

Example (2): A 385-gram chunk of iron is heated to 97.5°C. Then it is immersed in 247 grams of water originally at 20.7°C. When thermal equilibrium has been reached, the water and iron are both at 31.6°C. Calculate the specific heat of iron.


– The amount of heat gained by the water as it is warmed from 20.7°C to 31.6°C is the same as the amount of heat lost by the iron as it cools from 97.5°C to 31.6°C.

– We can equate these two amounts of heat and solve for the unknown specific heat.


Temperature change of water = 31.6°C – 20.7°C = 10.9°C
Temperature change of iron = 97.5°C – 31.6°C = 65.9°C

Heat and Temperature


  • Principles of Inorganic Chemistry / Brian W. Pfennig / 1st ed, 2015 /John Wiley & Sons, Inc/ USA.
  • Inorganic Chemistry /Peter Atkins, Tina Overton, Jonathan Rourkel, Mark Weller, Fraser Armstrong, Mike Hagerman / 6th ed, 2014 /W. H. Freeman and Company/ New York, USA.
  • Complete Chemistry for Cambridge IGCSE RG Student book/RoseMarie Gallagher, Paul Ingram / 3rd ed, 2014 / Oxford University Press / USA.

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