Unit 2: Essentials of Chemistry

1. Unit 2: Essentials of Chemistry Chapter 1-2, 4-5

2. Objectives 8 explain the nature of science including the use of the validity of the scientific method and the difference between a hypothesis, theory and law 9 explain the three major states of matter and their physical and chemical characteristics 10explain and give examples of physical properties and chemical properties 11explain and give examples of physical changes and chemical changes 12identify the five pieces of evidence that a chemical reaction has occurred 13define and classify matter including pure substances mixtures. 14 understand the historical progression of structure of the atom including such models created by Democritus, Dalton, Thomson, Rutherford, Bohr, and Schrodinger 15define and identify examples of the laws of conservation of mass, definite proportions, and multiple proportions 16identify the structure of an atom including the relative masses of a proton, neutron, and electron, their relative charges, and locations in the atom 17define and identify isotopes and ions 18define and distinguish between atomic number and mass number and identify the parts of the nuclear symbol 19define and calculate average atomic mass in amu’s 20understand and record the arrangement of electrons in an atom including Hund’s rule, the Pauli exclusion principle, and the Aufbauprincipleby writing the orbital notation and electon configuration for specific elements

3. 8 Explain the Nature of Science • The nature of science refers to how science is actually performed. • It explains the scientific view point on solving problems and why certain ideas are not considered. • There are several nature of science ideas which will be discussed on the next few slides

4. Nature of Science Ideas • While science strives to be objective, it cannot be entirely. • It is human nature to use previous experiences and knowledge. This allows scientists to make logical conclusions even though it is not completely objective. • For example: Imagine you were trying to determine an object in a shoebox. An objective scientist would have to try everything. This would be tedious and waste time because certain objects would not fit (like an elephant). By using prior knowledge, certain objects can be eliminated and the object can be identified faster.

5. Nature of Science Ideas • Science needs experimental evidence. • In order for something to be considered scientifically supported, it must be able to be tested. • Only testable evidence will be accepted by the scientific community. They will not accept unless some experiment supports it. Great scientists have proposed ideas well ahead of their time, but these ideas were not accepted because there was no evidence to support their claims.

6. Nature of Science Ideas • Science neither accepts or denies the existence of the supernatural. • Supernatural beings are considered to be outside the realm of science. • A supernatural being can change the natural world at their will so science cannot test for their existence. • Because science cannot test it, science cannot make a statement one way or the other on the existence of the supernatural. • Science only tries to develop conclusions based on ideas that can be tested.

7. Nature of Science Ideas • Science cannot prove anything 100% which makes it a flexible subject. • As that science relies on data and evidence, it can only support ideas. • As new evidence is generated, this support can strengthen or weaken. • Often new technology allows for new information to be discovered. • For example: In ancient Greece, it was believed atoms were the smallest particles. However, in the early 1900s, scientists used new technology to show that atoms were made of sub-atomic particles (protons and electrons). The idea that atoms were the smallest particles was no longer supported with this new data.

8. The Scientific Method • When looking at how science is performed, the scientific method must be discussed. • It is designed with a few steps which are: • Observe a situation • Make a hypothesis • Test the hypothesis • Make a conclusion • Other sources may provide more steps to the method but it can be condensed to these four.

9. The Scientific Method • The scientific method is an excellent problem solving guide because it is flexible. • It allows for the scientists to start at any point and proceed with solving the problem • For example, imagine you are performing an experiment and observe a change unrelated to your experiment. This could be the beginning of a new hypothesis yet it did not come from simply observing the natural world. • This method for solving problems works for all problems and not just scientific ones.

10. Hypothesis, Theory, and Law • These three terms are often confused when discussing the nature of science. • A hypothesis is an educated analysis of an observation. It requires testing and is essential an explanation of a phenomena. • A theory is a hypothesis with enough evidence to be accepted as fact. It explains the how and why qualitatively. • A law is a hypothesis with enough evidence to be accepted as fact. It describes the what quantitatively.

11. Hypothesis, Theory, and Law Hypothesis Theory Law It is possible to change a theory or a law if there is evidence to support such a change. Can become with enough experimental evidence Can become with enough experimental evidence Both are accepted as fact and cannot become the other

12. 9States of Matter • There are actually five known states of matter. Three are common and will be discussed in greater detail. • Solid • Liquid • Gas • Plasma (occurs in the sun) • Bose-Einstein Condensate (occurs when temperatures approach -273°C)

13. Solids • Solids have a defined shape and a defined volume. • The atoms that make up its structure are very close together and are limited in their movement.

14. Liquids • Liquids have a undefined shape but a defined volume. • The atoms that make up its structure are free to move around within its volume.

15. Gases • Gases have an undefined shape and an undefined volume. • Their atoms are free to move about the entire volume presented.

16. 10.Properties • Each substance has a set of properties (descriptions). These can fall into two categories: Physical and Chemical Physical properties describe or measure the object without changing it. Chemical properties describe the substance’s ability to undergo a change. Example: color, smell, mass Example: ability to rust, flamability

17. 11.Changes • Each substance also can undergo a change and the same two categories apply: Physical and Chemical Physical change is a change in which the substance is altered but can be returned to its original state. Chemical change is a change in which the substance is altered at the molecular level and cannot be returned to its original state. Example: smashing, tearing Example: baking, burning

18. Phase changes • Phase changes occur when one state of matter becomes another. • Melting S  L • Freezing L  S • Vaporization L  G • Condensation G  L • Sublimation S  G • Deposition G  S • Since these changes can be reversed to get the original back, phase changes are a type of physical change.

19. 12.Evidence for a chemical change • Chemical changes can be more challenging to determine. • To help in this matter, there a some pieces of evidence to look for in the event of a chemical change. • Color change • Formation of a Precipitant • Emission of heat • Emission of light • Emission of a gas The videos these are linked to are from youtube.com as of July 26, 2011

20. 13.Classifying Matter • As discussed earlier, there are three main types of matter: solid, liquid, and gas. • Matter can be broken down into two categories: Pure substances and Mixtures • A pure substance consists of only one component and it has unique chemical and physical properties. • A mixture is a combination of two or more pure substances.

21. Pure Substances • Pure substances can be broken down into a few different terms. All of the following have unique chemical and physical properties.

22. Allotropes • An allotrope is an unique type of molecule. • Allotropes are atoms of the same type that bond to themselves in multiple ways. • For example: Oxygen • Oxygen gas is O2 which means there are two oxygen atoms • Ozone gas is O3which means there are three oxygen atoms.

23. Mixtures • Mixtures are two or more pure substances and can be mixed differently. • Homogenous mixtures are thoroughly mixed and the parts are uniform throughout. • Any solution is a good example. • Heterogeneous mixtures are not uniform and the parts can be seen mixed throughout. • Most mixtures are heterogeneous and a good example of this would be a chocolate chip cookie.

24. 14.History of the atom

25. Dalton’s Atomic Theory • Elements are made of tiny particles called atoms. • Atoms of an element are different from other elements and can be distinguished by their atomic masses. • All atoms of a given element must be identical in properties. • Atoms of an element can combine with atoms of a different element to for compounds. • Chemical reactions rearrange the atoms but cannot create or destroy atoms. Return

26. Thomson’s Plum Pudding Model • As that Thomson discovered the electron, that meant the atom contained smaller parts. • This would change the model used for the atom. • The model used to be a solid sphere • The plum pudding model used electrons as the “plums” and the rest of the atom as the “pudding”. • The plums were negative and the pudding was positive. Return

27. Gold Foil Experiment • The gold foil experiment was conducted by Ernest Rutherford and his graduate students, Hans Geiger and Robert Marsden. • By shooting alpha particles at a gold foil, they noticed the particles essentially went straight through. • This led them to conclude the atom was mostly empty space with a dense positive core (protons. Return Picture was taken from: http://www.kentchemistry.com/links/AtomicStructure/rutherfordtutorial.htm on July 28, 2011.

28. Bohr Model of the Atom • Niels Bohr created an atomic model after doing work with the color spectra emitted from a hydrogen atom. • His model is sometimes called the solar system model. • The following link provides a more detailed description (the first 15 slides covers Bohr): http://science.sbcc.edu/physics/solar/sciencesegment/bohratom.swf Return

29. Schrodinger’s Orbitals • The current model of the atom uses the orbitals discovered by Schrodinger. • Within each energy level, there exists four kinds of orbitals: s, p, d, and f. • Each can hold a certain number of electrons. • The shape of each orbital is shown on the next slide. • The image was taken from: http://chemwiki.ucdavis.edu/Physical_Chemistry/Quantum_Mechanics/Atomic_Theory/Electrons_in_Atoms/Electronic_Orbitals on July 28th, 2011.

30. Return

31. 15.Define and Identify Three Laws • Law of Conservation of Mass • The law of conservation of mass states that mass can be mass can neither be created or destroyed. • This means that the products of a chemical reaction will have the same mass as the reactants.

32. 15.Define and Identify Three Laws • Law of Definite Proportions • This law states that any sample of a compound has the same composition. • This means that water will always be H2O whether it is found in Iowa or somewhere else.

33. 15.Define and Identify Three Laws • Law of Multiple Proportions • This law states that the mass ratio for one of the elements in a compound that combines with a fixed mass of another element can be expressed in small whole numbers.

34. 16.Identify the structure of the atom • The atom is constructed of three basic subatomic particles. • In the last 20 years, it has been discovered that quarks make up both protons and neutrons.

35. Structure of the Atom-Bohr Model This picture came from http://www.universetoday.com/82128/parts-of-an-atom/ on August 31, 2011.

36. 17.Isotopes and Ions • Each atom can have different variations. • All atoms are identified by the number of protons they contain. • Oxygen will always have 8 protons. If you added an additional proton, the atom would no longer be oxygen (it would be fluorine). • The number of neutrons and electrons can vary slightly.

37. Isotopes-changes in neutrons • Protons contain a positive charge. When an atom becomes large, it contains several protons. That much positive charge in one location is unstable. • Neutrons, which have no charge, act as spacers in between the protons. • An isotope is the name of an atom with different amounts of neutrons. • Oxygen for example has three common isotopes. • Oxygen-16 has 8 neutrons • Oxygen-17 has 9 neutrons • Oxygen-18 has 10 neutrons

38. Ions-changes in electrons • Electrons fill the orbitals surrounding the nucleus. • This is known as the electron cloud. • In a neutral atom, each proton has an electron in the electron cloud. • When an atom becomes charged, the number of electrons no longer matches the protons. • Positive charges indicate a loss of an electron. • Negative charges indicate a gain of an electron.

39. 18.Nuclear Symbols • Nuclear symbols are useful for determining the amount of electrons, protons, and neutrons in an element. Mass number: Shows the number of neutrons and the number of protons. Charge: Indicates the difference between the electrons and protons. Negative charges indicate more electrons while positive charges indicate more protons. Atomic Number: Shows the number of protons.

40. 19.Average atomic mass • On the Periodic Table, the mass listed is the average atomic mass. • This is an average of all the naturally occurring isotopes of an atom. • Calculating an average for a large amount of particles can be challenging.

41. Average Atomic Mass • Carbon has three common isotopes. • Carbon-12 • Carbon-13 • Carbon-14 • These three common isotopes do not come in equal amounts though: • 98.89% is carbon-12 • 1.10% is carbon-13 • 0.01% is carbon-14

42. Average atomic mass • To determine the average atomic mass, the following formula should be used: (Atomic Mass x Percent Abundance) + (Atomic Mass x Percent Abundance) +…..=average atomic mass • So in the case of carbon: (12.00 amu x .9889) + (13.00 amu x 0.0111) + (14.00 amu x 0.0001) = 12.01 amu

43. 20.Arrangement of Electrons • According to Bohr and Schrodinger, electrons surround the atom in energy levels that take the form of orbitals. • The Periodic Table is broken into four sections that correlate to the orbitals • The placement of electrons within these orbitals follow set rules.

44. Arrangement of Electrons • Hund’s Rule • Electrons will fill the available orbitals at a certain energy level before pairing. • Pauli Exclusion Principle • Only two electrons can occupy each orbital and their spins will be opposite. • Aufbau Principle • The lowest-energy orbitals will be filled first.

45. Below is a depiction of how the Periodic Table shows the orbitals (SPDF). S P F D

46. Energies of Each Level • The orbitals correspond to different levels of energies. • In general, each row on the Periodic Table represents a different level of energy. • It gets more complicated farther down the Periodic Table.

47. Energies of Each Level • 7p ___ ___ ___ • 6d ___ ___ ___ ___ ___ • 5f ___ ___ ___ ___ ___ ___ ___ • 7s ___ • 6p ___ ___ ___ • 5d ___ ___ ___ ___ ___ • 4f ___ ___ ___ ___ ___ ___ ___ • 6s ___ • 5p ___ ___ ___ • 4d ___ ___ ___ ___ ___ • 5s ___ • 4p ___ • 3d ___ ___ ___ ___ ___ • 4s ___ • 3p ___ ___ ___ • 3s ___ • 2p ___ ___ ___ • 2s ___ • 1s ___ As you progress from the 1s to the 7p, you increase the amount of energy. Notice how the d-sublevel is always after the s-sublevel of the previous energy level (example 3d follows 4s). Notice the f-sublevel follows the s-sublevel of two energy levels before it (example: 4f follows 6s). Energy

48. 2.14 Orbital Notation • To show orbital notation, the three rules must be followed. • Therefore, start at the lowest energy level. • Designate an electron by drawing an arrow • The arrow indicates the spin • Place one electron in each orbital until they each have one on that level. Then go back and pair them. • Only two electrons fit in each orbital. • The arrows must point in opposite directions to show opposite spins.

49. Orbital Notation • Look at the element sulfur-32. • First determine the number of electrons • Since sulfur has an atomic number of 16, it has 16 protons, 16 electrons, and 16 neutrons. 3p ___ ___ ___ 3s ___ 2p ___ ___ ___ 2s ___ 1s ___ Place electrons in the first energy level and continue up. When 2p is hit, fill each orbital and then go back and pair. Finish by repeating the process used for 2p but stop when you hit 16 arrows. This is the orbital notation.

50. Electron Configuration • Orbital notation can take up a lot of space. • It does a nice job of giving a visual of the location of each electron. • Electron configuration is a shorthand notation for determining the location of the electrons. • It follows the same rules but is slightly easier to write.