Regents Chemistry & The Scientific Method – NYS Chemistry – The Source for Regents Chemistry Review (2023)

1.Regents Chemistry is the study of matter and the changes that matter undergoes.

Pure chemistry is chemistry done for the sake of gaining new information. An example would be any type of research, such as an analytical chemist determining the composition of an unknown compound, or a biochemist determining the chemical structure of a newly discovered virus.

Applied chemistry is chemistry done to improve our lives. It does not discover new information about chemistry, but makes practical use of existing chemical knowledge. An example would be manufacturing rayon or nylon (synthetic fibers) and producing clothes.
2. The five divisions of chemistry are

a. organic chemistry – the study of carbon based compounds – includes plastics, rubbers, petroleum products and many synthetic fibers – rayon, nylon etc.

b. inorganic chemistry – the study of compounds which do not contain carbon – acids, bases, salts etc.

c. analytical chemistry – studies the composition of substances – used extensively in research and forensics

d. physical chemistry – determines and describe the physical structure of matter to explain why it behaves the way it does

e. biochemistry – studies the composition and the changes in composition of the matter which makes up living things

3. Two basic assumptions that make the endeavor of science possible are

a. the world is uniform – the laws of nature are consistent over space and time

b. humans have the capacity for understanding the world

These may seem rather obvious to us, living in a “scientific era,” but to a prescientific mind set it was not obvious. In a pagan or superstitious world view matter was controlled or, at least affected by demons and spirits who changed things according to their whims (whether mischievous or benevolent).

4. The three major activities the science engages in are

a. Observation – gathering information about the world -either directly or indirectly through our senses

b. Generalization – making broad statements about reality based on observations (hypotheses)

c. Verification – Testing of generalizations to see if they hold true – usually by use of controlled experiments. Controlled experiments have only a single variable (if possible), a control group (not exposed to the variable – used for comparison), and an experimental group – exposed to the variable).

Consider the story of Sir Isaac Newton. An apple hits Sir Newton while sitting under the apple tree (observation). Sir Newton reflects on this observation and others and comes to a broad understanding that all matter seems to be attracted to each other (generalization). Sir Newton then tests this generalization by observing and measuring matter to see if all matter does indeed attract all other matter. When Newton collects enough data and uses his generalization to make predictions of the behavior of matter which hold true he can conclude that his generalization is true and states his generalization as a theory or law (in this case the law of gravity).

5. There are many forms of the scientific method. All involve the steps above. One form of the scientific method is below :

a. State the problem

b. Make observations (collect data – factual observations)

c. Form a hypothesis- educated guess as to the solution of the problem

d. Test the hypothesis

e. Form a conclusion (e.g. “Hypothesis is true” or “Hypothesis is false.”

f. Repeat the experiment to be sure results are consistent.

6. If the results are consistent the conclusion may eventually result in a theory (scientific statement which explains a natural phenomenon), a law (a description {often mathematical} of a natural phenomenon) or a model (a possible explanation or description of a natural phenomenon). Models result when the hypotheses can not be fully tested. Models make sense of the data, but cannot be verified experimentally.

7. Scientific knowledge is:

a. tentative – always subject to change based on new observations (it was once “scientifically true” that the world was flat and that the earth was the center of the solar system

b. limited – science cannot study all things – the existence of God, the perception of beauty etc.

c. based on observation

d. a blend of logic and imagination (especially models)

e. explanatory and predictive – theories are used to explain and predict the behavior of matter

f. the result of human activity and will have social, ethical and moral implications (e.g. bombs, birth control pills) and economic implications (e.g. new industries).

8. Matter – anything which takes up space and has mass

Mass – a measure of the amount of matter an object has

Substance – a particular type of matter which has a uniform and definite composition (e.g. sodium (Na), water (H2O).

8a. Intensive properties are properties that do not depend on the amount of the substance present.

An example of an intensive property is density. Whether you have a small piece of lead, or a large piece of lead, its density is always 11.4 g/cm3.

Extensive properties do depend on the amount of the substance present. An example of an extensive property is volume. Different sized pieces of lead will have different volumes.

9. Physical properties are those properties which can be observed without the substance changing into a new substance. (e.g. brittleness of chalk, color of the sky)

Chemical properties are those properties which cannot be observed without the substance changing into a new substance. (e.g. Wood burning)

10. Four states of matter :

• solid – has a definite shape and volume
• liquid – has a definite volume, but not a definite shape (takes the shape of its container)
• gas – has neither a definite shape or volume
• plasma – highly energized state of matter in which the electrons have separated from the nuclei creating a mixture of electrons and positive ions (e.g. the sun is made of plasma, only here the matter is so ionized it is almost pure nuclei -one tablespoon has an estimated mass of about 2 tons).

11. Gases are substances which are in the gaseous state under normal environmental conditions. Vapors are the gaseous states of substances which are normally solids or liquids.

12. Physical changes are changes in the physical properties of a substance (do not result in a new substance) – ex. – breaking a pencil (changing size).

Chemical changes are changes which result in a new substance. For example, burning wood changes it into mostly water and carbon dioxide.

13. A mixture is a physical combination of two or more substances. Homogeneous mixtures have a uniform composition (e.g. salt water), heterogeneous mixtures are not uniform in composition (a bucket of dirt).

14. A solution is a homogeneous mixture where one substance (the solute) is completely and evenly dispersed in another substance (the solvent). Salt dissolving in water is a physical change – the salt is still salt and the water is still water. A phase is any part of a system with uniform composition and properties. A homogeneous mixture, such as a solution is made up of only one phase. A heterogeneous mixture consists of more than one phase.

15. Distillation is a process in which a liquid mixture is boiled to produce a vapor that is condensed back into a liquid. This can remove the solids in water, such as tap water. This process can also separate mixtures of two or more liquids which have different boiling points.

16. The two groups of substances are

a. elements – simplest forms of matter which cannot be broken down into simpler substances by ordinary chemical means. Elements are made up of only one kind of atom

b. compounds – substances made up of two or more elements chemically combined, can be broken down into simpler substances by ordinary chemical means

17. See the periodic table.

18. Energy is the ability to do work. The two categories are :

a. potential energy – energy that is stored, or energy of position
b. kinetic energy – energy of motion

19. Six forms of energy :

a. chemical energy – (potential) the energy stored in the chemical bonds of a substance
b. nuclear energy – (potential) the energy stored in the nuclei of matter (released by fission or fusion according to Einstein’s equation E = mc2, where E= energy, m = mass, c = speed of light)
c. electrical energy – (kinetic) the energy of moving electrons
d. mechanical energy – (kinetic) the energy of a moving object
e. radiant energy – (kinetic) the energy of traveling waves (light, sound, x-ray etc.)
f. heat energy – (kinetic) the energy of the moving particles of a substance

20. Law of Conservation of Energy (First law of Thermodynamics) – energy can neither be created nor destroyed (The amount of energy in the universe is constant).

21. Second Law of Thermodynamics – in any spontaneous change useful energy is lost -although the amount of energy in the universe is constant, useful energy is decreasing. Eventually the universe will die a “heat death” – all useful energy will be lost and the universe will be a uniform temperature.

22. A chemical reaction is a reaction in which a chemical change takes place – one or more substances become new substances. A physical reaction results in the changes of the physical properties of one or more substances, but no new substances are produced.

23. Reactants are the substances that are changed in a chemical reaction and the products are the substances produced in a chemical reaction. For example, in the summary chemical reaction of photosynthesis, water and carbon dioxide are the reactants and glucose and oxygen are the products.

6 CO2 + 6 H2O —> C6H12O6 + 6 O2
reactants products

24. Factors which indicate a chemical change include :

– a change in color or odor
– a change in temperature – energy is absorbed or released
– a change in state (e.g. gas being formed from a liquid, or a solid being produced from a mixture of two liquids)

25. A precipitate is an insoluble solid produced by a chemical reaction.

26. The Law of Conservation of Mass states that in any physical or chemical reaction mass is conserved – neither gained nor lost.

27. Measurement is important to science because measurement gives meaning to observations. Consider the statement “That reaction produced a lot of energy.” To a biochemist it may mean a ten degree rise in temperature, to a nuclear physicist it may mean the release of enough energy to obliterate a mountain!

Quantitative measurements (think of the word “quantity”) result in a definite form, usually involving numbers.

Qualitative measurements are descriptive and are nonnumeric.

28. Accuracy refers to how close a quantitative measurement is to the accepted value and precision is how repeatable the results of a quantitative measurement are. Lets say that I know an object has a mass of 5.00 grams. I have two sets of five students measure the mass of the object giving the following results :

Set 1 Set 2
student 1 : 5.00 g student 1 : 6.50 g
student 2 : 3.50 g student 2 : 6.49 g
student 3 : 8.20 g student 3 : 6.51 g
student 4 : 6.50 g student 4 : 6.50 g
student 5 : 5.78 g student 5 : 6.50 g

Only student 1 in data set 1 is accurate and the set of data shows little precision. In data set 2 no one is accurate (remember the accepted value is 5.00g), but the data set shows good precision. Imagine the following targets taken from a firing range.

Sample ImageSample ImageSample Image
Target 1 shows neither accuracy nor precision. Target 2 shows good precision, but poor accuracy. Target 3 shows both good accuracy and good precision.

29. Error is a measure of the accuracy of a measurement. Absolute error (Ea) is the absolute (always positive) difference between the accepted value and the observed (measured) value.

Absolute error = observed value – accepted value (Ea = O-A)

Relative error is the absolute error expressed as a percentage of the accepted value.

Relative error (Er) = Ea/A x 100%

Relative error is a better expression of the accuracy of a measurement. An absolute error of .005 grams may sound great, but if the accepted value is .0001 g, the Er is 5,000 %!

30. Deviation is a measure of the precision of a set of measurements. Put in another way, it is a measure of the uncertainty of a set of measurements. Absolute deviation (Da) is the difference between each measurement in a set of measurements and the average of the set of measurements.

Absolute deviation = absolute value of the observed value – average of observed values (Da = |O-Davge| )

Consider the following set of data :

Data Da
5.00 g .06 g
5.06 g .00 g
5.08 g .02 g
5.10 g .04 g
avge = 5.06 g avge = .03 g

The average of the four measurements is 5.06 g and the Da is calculated by the above equation. How should the final answer be expressed? The best estimate is the average of all the measurements, but can the person measuring honestly say that it is exactly 5.06 g? No, there is some uncertainty in the answer and it should be indicated in the expression of the answer. The best possible answer is expressed as the average of the measurements plus or minus the average of the deviations.

Answer = average + average of the deviations

In our example this is 5.06 + .03 g. This means that the best answer is 5.06 grams, but it really might be .03 grams more or .03 g less.

31. Significant figures are all those digits in a measurement which can be known precisely and a last digit which can be estimated. Significant figures are those digits in a measurement which are actually measured and are an indication of the precision of the instrument used to make a measurement. A measurement of 1 meter has only one significant figure and has little precision. A measurement of 1.000 meters has four significant figures and has a much greater degree of precision. Consider the measurement at the arrow on the following centimeter ruler.

Sample Image

A measurement made at the arrow would be 1.94 cm. The digits 1 and 9 are known because we can see the lines which indicate these numbers. The last digit, 4, is estimated. It would not be correct to round the measurement up to 2.0 or down to 1.9, because we can see that it is somewhere in between. The ruler has a degree of precision which is to the .01 (hundredths) cm place. What should the measurement indicated on the following ruler be?

Sample Image

The measurement should be 3.00 cm. This ruler has lines at the tenths of a cm place and therefore the measurement can be carried one place further – the last estimated digit. Contrast this to the ruler below.

Sample Image
The measurement should be 3.00 cm. This ruler has lines at the tenths of a cm place and therefore the measurement can be carried one place further – the last estimated digit. Contrast this to the ruler below.

A measurement made at the arrow would be 1.5 cm. This ruler has a lesser degree of precision.

32. Significant figures are used in measurements -where instruments with certain degrees of precision are used. Significant figures are not used in defined quantities. For example, by definition 1 meter equals 1 000 millimeters. I do not need to write 1.000000…. m, because a meter is always exactly a meter, it has an unlimited number of significant figures.

33. The following are the rules for determining which digits in a number are significant and which are not.

• All nonzero digits are significant.

• Zeroes appearing between nonzero digits are significant.

• Zeroes appearing in front of all nonzero digits are not significant (e.g. the zeroes in the number 0.0056 are not significant).

• Zeroes to the right of a nonzero digit and to the right of a decimal place are significant. (e.g. 2.000 has 4 sig figs, while 0.0200 only has three sig figs).

• Zeroes to the right of nonzero digits and to the left of a decimal are confusing. For example, if I count 30 students, the last zero is significant because it is a measured quantity. I have counted my students with a degree of precision to the one’s place. If I look at a crowd and estimate that there must be 500 people, the zero in the one’s place is not significant. The number does not have a degree of precision to the one’s place. To avoid this confusion numbers should be written in exponential, or scientific, notation (see obj. 35).

• It is extremely important that you get a good handle on significant figures.

34. In addition and subtraction the sum or the difference can have nor more decimal places than the number in the calculation which has the least number of decimal places. For example, if I subtract 1.2 from 3.40, the final answer is 2.2, not 2.20. If I add 2.20 and 3.1, the final answer is 5.3.

In multiplication and division the answer can have no more significant figures than the number in the calculation with the least number of significant figures. For example if I multiply 2.6 and 4.555, the final answer can only have two significant figures.Regents Chemistry & The Scientific Method – NYS Chemistry – The Source for Regents Chemistry Review (1) Regents Chemistry & The Scientific Method – NYS Chemistry – The Source for Regents Chemistry Review (2) Regents Chemistry & The Scientific Method – NYS Chemistry – The Source for Regents Chemistry Review (3) Regents Chemistry & The Scientific Method – NYS Chemistry – The Source for Regents Chemistry Review (4)

Note the difference between addition-subtraction and multiplication-division. In addition or subtraction it is the number of decimal places that matter, in multiplication and division it is the total number of significant figures that matter.

The reason for these rules is that measurements must not be presented to a greater degree of precision than is actually known. For example if two measurements are to be added, 3.2 cm and 4.0001 cm the sum would be 7.1001 cm. However the first measurement (3.2 cm) is not very precise, it is known only to the tenths of a cm place. The hundredths, thousandths and ten thousandths place are not known and should therefore not be reported in the sum.

36. The system of units used for measuring in science is the SI (after the French “Le Systèmes International d’Unités). It is commonly referred to as the metric system, but is actually a revised version of the metric system. The advantages of this system is that it is used worldwide, making it easier for scientists to compare and share data, and that it is based on the decimal system, making it easy to convert from one unit to another.

37. The basic unit of length is the meter. The basic unit of volume is the cubic meter (m3). The basic metric unit of mass is the kilogram.

38. Volume is the amount of space an object occupies. Mass is the amount of matter that an object has. Weight is the pull of gravity on an object.

39. Metric prefix Symbol Meaning
kilo- k 1000
deci- d .1 (one-tenth)
centi- c .01 (one-hundredth)
milli- m .001 (one-thousandth)
micro- µ 10-6 (one-millionth)
nano- n 10-9 (one-billionth)

40. A liter is the most commonly used metric unit of liquid volume and is equal to 1000 cm3 (a 10 cm x 10cm x 10 cm cube). A kilogram is the basic metric unit of mass and it is equal to 1000 grams (a gram is so small that it is too inconvenient to use as the basic metric unit of mass).

41. Density is the ratio of the mass of an object to its volume. Density is calculated by dividing the mass of an object by its volume. For an example, if a 10.0 kg object occupies a volume of 2.0L its density is 10.0kg/2.0L which is 5.0 kg/L. Note the complex units which indicate the ratio of mass to volume- kilograms per liter.

42. Specific gravity is the ratio of the density of a substance to the density of a reference substance which is usually water. Specific gravity is calculated by dividing the density of the substance in question by the density of the reference substance. For example, gold has a density of 19.3 g/cm3 at 25ºC. Using water as a reference substance (density = 0.997 g/cm3 at 25ºC) we can calculate the specific gravity of gold by dividing 19.3 g/cm3 by 0.997 g/cm3 to get 19.4 as the answer. Note units cancel out and therefore specific gravity does not have any units.

43. A hydrometer is the device used to measure the specific gravity of liquid substances. (See fig. 3-17 p. 72 of text).

44. Temperature is a measurement of the average kinetic energy of the particles of a substance. Five different devices used to measure the temperature are :

• bimetallic thermometer – consists of two strips of two different metals which expand at different rates to indicate temperature

• alcohol or mercury filled thermometer – the liquid is contained in an evacuated glass rod. the liquid expands and contract to indicate the temperature on a scale

• thermistor – made up of a metal oxide with a current running through it. the conductivity of the metal oxide changes with changes in temperature which can then be measured to indicate temperature

• liquid crystal display – contain a liquid crystal which changes color to indicate the temperature

• thermocouple – consists of two different metal wires joined at two points. One junction is used as a reference at a known temperature. The other junction is put at the place where the temperature is to be measured. the differences in temperature create an electrical flow which can be measured and used to calculate temperature.

45. The Celsius scale of temperature has the freezing point of water at 0 degrees and the boiling point of water (at 1 atmosphere of pressure) at 100 degrees. It is used primarily by the sciences other than physics. The Kelvin scale is based on absolute zero, the theoretically lowest possible temperature where all molecular motion would cease. It is used primarily by physicists. We will use it when working with the gas laws. Note there are no negative values on the Kelvin scale- that’s why it is absolute zero!

46. The Kelvin scale is 273 higher than the Celsius scale. the following equations can be used to convert between these scales.

K = ºC + 273 and ºC = K-273


The following is a list of equations you will need to be able to use.

1. Ea = IO-AI
2. Er = Ea/A x 100%
3. Da = IO-DavgeI
4. D = m/V
5. specific gravity of x = density of x/density of water
6. K = ºC + 273

Where : Ea = absolute error
Er = relative error
A = accepted value
Da= absolute deviation
Davge= average deviation
O = observed measurement
D = density
m = mass
V = volume

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