Unit 1: Matter- Qualitative & Quantitative Properties
An Overview of Chemistry
Fundamentally, chemistry can be summarized with two concepts:
- Matter is composed of various types of atoms
- One substance changes into another by reorganizing the way atoms are attached to one another.
The various properties of matter are determined by the properties of the atoms and/or molecules that comprise it. Chemistry is the study of these properties.
The Scientific Method
In chemistry, we use the scientific method to observe the macroscopic effects of chemical changes.
- Make observations and collect data
- Suggest a possible explanation for this data (hypothesis)
- Do an experiment to TEST your possible explanation (repeat as needed)
- Create a theory to describe repeatable results of your experiment.
- Theories are just possible explanations for why a phenomena is occurring
- Scientific laws are a summary of an observed natural behavior
- A law summarizes what happens while a theory is an attempt to explain why it happens
Examples:
Atomic Theory- Dalton proposed the atom to be a small indivisible sphere of mass
Law of Conservation of Mass- In a chemical reaction, mass is neither created nor destroyed.
Classifying Matter
The simplest way to classify matter is by state:
Solids, Liquids, and Gases all have unique properties
You need to remember the basicsfrom Chem 1:
A second common way to classify matter is by its composition:
A pure substance is made up of only one component and its composition is invariant. Can NOT be separated by physical means.
A mixture, by contrast, is a substance composed of two or more components in proportions that can vary from one sample to another. Mixtures CAN be separated by physical means.
Separation of Mixtures
Many times chemists need to separate mixtures into their components. (Think of refineries that purify various components of crude oil.) We use the fact that each component of a mixture has its own unique chemical and physical properties to create a method for separating the substances. You MUST know these various techniques:
Decanting- separation by different states (solid vs liquid)
Filtration- separation by size (small vs large)
Distillation- separation by boiling point (volatile liquid vs. nonvolatile liquid)
Physical vs. Chemical (More review from Chem 1)
Physical Properties- properties displayed without changing composition
Include odor, taste, color, appearance, normal melting point, normal boiling point, density
Chemical Properties- only displayed by a chemical reaction
Include corrosiveness, flammability, acidity, toxicity
Physical Change- substance is not changed at the atomic or molecular level
Include cutting, tearing, dissolving, breaking, boiling, melting, sublimation
Chemical Change- atoms are rearranged to make one or more new substance
Any chemical reaction or process
Law of Conservation of Energy
-Energy is always conserved in a physical or chemical change; it is neither created nor destroyed.
-Systems with high potential energy tend to change in a direction that lowers their potential energy, releasing energy into the surroundings.
Measurement
- Quantitative measurements are numerical data (length, temperature, mass, density, volume)
- When observing mass, be sure to make the distinction between the amount of matter (mass) versus the pull of gravity of that mass (weight)
- Qualitative measurements are taken with the senses (odor, color, sound, feel, taste)
- All measurements should be taken using the International System, or SI System. You should be familiar with units of this system (Liter, Kilogram, Meter, Celsius)
- Units of this system can be modified with prefixes to change their meaning. Be sure to know these prefixes and how to use them.
Temperature
There are three systems for measuring temperature. We focus on Kelvin and Celsius. Be sure you know how to convert between all three systems.
Example: Convert 212oF to Kelvin
Density and Classification of Matter
Density is a derived unit (a combination of two other units, volume and mass)
We can measure, classify and convert matter.
- Density is a property of all pure samples of matter. It measures mass per unit of volume.
- D = m/v
Practice: Determine the density of the following substances in g/ml.
- 500 g of Mercury takes up a volume of 50 ml
- Pure water has a density of 1g per ml. Calculate the density of salt water if 1L of this solution weights 1100 grams.
Put the three states of matter in order from most to least dense. Include particle diagrams.
Intensive property- property that is independent of the amount of the substance. These are what we use to identify unknown substances.
Example: Density, Normal Boiling Point, Reactivity
Extensive property-property that depends on the amount of the substance.
Example: Mass, Volume, Length
Uncertainty in Measurement
Uncertainty results from LAB EQUIPMENT and should ALWAYS be taken into account.
Measurements should include ALL CERTAIN
digits plus ONE UNCERTAIN digit.
Ex. Reading a thermometer, buret, balance, etc.
- Read the above rulers to the correct number of significant digits.
A:______B: ______
Significant Figures and Calculations
When taking measurements, the digits present in the measurement that were both read and estimated are called significant digits.
Sig Fig Cheat Sheet: (See Pg. 22 in Tro for all Sig Fig Rules)
Leading Zeroes are NEVER significant
Zeroes between two nonzero numbers are ALWAYS significant
Trailing Zeroes are only significant IF there is a decimal in the number. If there is no decimal in the number, trailing zeroes are ambiguous.
Examples: Write in the number of significant digits.
______7.0023______4.00______0.03000 ______0.009______1,000
Significant figures in calculations
When adding and subtracting, use the lowest number of decimals.
3.0 + 6.00 = ______
When multiplying and dividing, use the lowest number of total significant figures.
3.0 x 6.00 = ______
Exact numbers have unlimited significant figures and should not be used to determine the significant figures in a calculation
Example: There are EXACTLY 1000 mL in EXACTLY 1 L.
Accuracy & Precision
Any measurement is only as good as the instrument used to measure it. Scientists often repeat measurements several time to increase confidence in their results.
Accuracy- how close the measured value is to the actual value.
Precision- how close a series of measurements are to one another or how reproducible the measurements are.
- Measurements that are close to or the same as the accepted (“actual”) value are ACCURATE
- Measurements that are repeatable are PRECISE
- It is possible for a set of measurements to be BOTH precise and accurate or NEITHER.
- When measurements are not precise, they are said to be random.
- When measurements are precise, but not accurate, they are said to have systematic error.
Practice: Determine whether each set of data below is accurate, precise, both or neither.
(Given: The actual value is 50.0 mL)
Student AStudent BStudent CStudent D
Trial 1 50.1 ml 60.5 ml 75.0 ml 100.0 ml
Trial 2 51.2 ml 40.2 ml 74.5 ml 125.2 ml
Trial 3 49.8 ml 53.2 ml 75.2 ml 180.7 ml
Student A:Student B:Student C: Student D:
Practice: Determine whether each set of data below is accurate, precise, both or neither.
(Given: The true mass is 10.0 g)
Conversion between Units (Your Need To Know has all common metric conversions)
To convert between units, we utilize conversion factors.
Ex: There are 60 seconds in 1 minute There are 1000 milliliters in I liter
OR OR
Practice: Write conversion factors for
a. feet to inchesb. centimeters to meters
c. Years to seconds
Dimensional Analysis is the process you MUST use to convert in this class.
***You MUST show your work for full credit on any problem that requires conversion between units.
Practice:
- Convert 7 ft to inches
- Convert 855 cm to meters
- Convert 7.56 km to mm
- Convert 50 years to seconds
Units Raised to a Power
What is the area of a room that is 10 square meters (m2) in square centimeters (cm2)?
Calculate the displacement of a 5.70 L automobile engine in cubic inches (in3). Given: 2.54 cm = 1 in
Note: we can use density as a conversion factor!
Practice: Convert 1500 g of salt solution to Liters of solution (Given: density = 1.2 g/mL).
The Early History of Chemistry
Greeks- in 400 BC, the Greek Aristotle determined that there were four fundamental substances (earth, air, fire and water).
- Although the greeks thought that matter was composed of particles, they were unable to test these theories due to lack of materials.
- Democritus- stated that matter was composed of small, indivisible particles called “atomos”
-tom = to divideatom = can not be divided
Alchemists- Over the next 2000 years, Aristotle’s works influenced a branch of science called alchemy that was focused on trying to turn cheap materials into gold.
- This led to the discovery of many elements and compounds like HNO3, H2SO4, HCl, and aqua regia (HNO3 and HCl). Many important scientific procedures were developed as well.
Robert Boyle- 1600’s
- The first person to perform quantitative experiments with a focus on gases
- Wrote a book called The Skeptical Chymist in which he described his theory that a substance was an element unless it could be broken down into simpler substances.
- His research destroyed the notion of only four elements.
Phlogiston theory (John Becher and George Stahl)
- Postulated that a substance called phlogiston flowed out of burning material during combustion and that materials stopped burning when the phlogiston level became too high.
- Joseph Priestly- 1700s- credited with the discovery of oxygen (deplogisticated air)
Fundamental Chemical Laws
Antoine Lavoisier
- He determined the nature of combustion with careful quantitative procedures
- He developed the Law of Conservation of Mass (mass is neither created nor destroyed)
- Disproved the phlogiston theory and named oxygen
- Wrote first chemistry textbook (Elementary Treatise on Chemistry)
Law of Conservation of Mass- Mass is neither created nor destroyed in a chemical reaction.
The same number and types of atoms will be found in the products as in the reactants. They are simply ______.
Proust- 1800s
- Developed the Law of Definite Proportion- a given compound always contains exactly the same proportion of elements by weight. If these amounts change, the identity of the compound changes.
- Stated that matter is discrete, or particulate, not continuous (atoms!)
John Dalton- 1800s
- Created the Law of Multiple Proportions based on the Law of Definite Proportions
- It stated that when two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers.
Law of Multiple Proportions Demonstrated with Oxygen and Nitrogen:
Modern View of Atomic Structure
*We will revisit the History of the Atom in our Atomic Structure unit, but here we will look at the basics
Nucleus (protons and neutrons)- very tiny
Electrons- determine chemical behavior
Isotope- atom with the same number of protons but different numbers of neutrons
Atomic number = number of protons
Mass number = number of protons + number of neutrons
Number of protons = number of electrons in a NEUTRAL atom
Counting by Weighing
If we were to takes the individual masses of five nails, each nail would probably have a small deviation in mass, but would be relatively the same (let’s say an average mass of 2.5 g per nail). If someone needed 1,000 nails for a building project, it is easier to take the mass of 1,000 nails (2500 g) than to mass them individually. Objects do not need to have identical masses to be counted by weighing; they behave as though they all have an identical mass- the average mass.
Just like these nails, atoms of the same element can have different masses. We call these isotopes.
These differences in mass are due to differences in the number of neutrons in the nucleus
(remember: the nucleus accounts for the vast majority of an atom’s mass)
Average Atomic Mass- an average of each isotope of an element, based on % abundance
To calculate:
% Isotope A (Mass of A) + % Isotope B (Mass of B) + … = Average Atomic Mass
Example: Find the average atomic mass of Element X if 1.40% of the sample is an isotope with a mass of 203.973 amu, 24.10% is an isotope with a mass of 205.9745 amu, 22.10% is an isotope with a mass of 206.9759 amu, and 52.40% is an isotope with a mass of 207.9766 amu.
Mass Spectroscopy
How do we know these isotopes exist? Mass Spectroscopy!
The mass spectrometer heats a sample of atoms until they vaporize. Then the atoms are passed through a beam of high speed electrons which knock off some of the atom’s electrons, creating positively charged ions. These ions are then thrown through an electromagnet. This field causes the path of the ions to curve. This deflection is dependent on the particle size (mass).
Look at the Mass Spec Data for Neon to the left. There are three peaks or bumps on the graph. Each peak represents the amount of particles of that particular mass that were detected. This shows us that there are three isotopes of neon in the sample.
Practice. Use the mass spectrum shown below to calculate the average atomic mass of this element, then identify the element.
The mass spectrum for ______The mass spectrum for ______
The overall mass spectrum for ______:
The Mole
The mole is the chemist’s way of counting atoms. The mole (Avogadro’s number) is equal to 6.022 x 1023 atoms of any element. The weight of one mole of a substance is equal to the average atomic mass on the periodic table.
The mole is standardized as the number of carbon atoms in 12.0 g of pure Carbon-12.
The mass of one mole of an element is equal to its atomic mass in grams.
Dimensional Analysis is the process you MUST use to convert in this class.
***You MUST show your work for full credit on any problem that requires conversion between units.
Example. Cody found a gold nugget that had a mass of 1.250 grams. How many moles was this? How many atoms?
Example: How many grams of calcium nitrate contain 24 grams of oxygen atoms?
Molar Mass- Mass (in grams) of one mole of a substance. (Also called molecular weight)
Example: Calculate the molecular weight of cisplantin (Pt(NH3)2Cl2)
Example: How many grams are in 3.25 moles of cisplantin?