Chapter 3: Atoms & moles

1. Chapter 3: Atoms & moles

2. Matter Anything that has mass and takes up space. Pure Substances Mixtures • 2 or more substances that can be separated by physical means Elements Compounds Heterogeneous Homogeneous • Chemical combination of 2 or more elements • Can be broken down by chemical means • Found on the Periodic Table • Cannot be broken down by chemical means into simpler substances • Mixed uniformly • Ex: Sugar Water, Salt Water • Not uniform • Ex: Salad, Pepperoni Pizza, Lucky Charms

3. Section 1: Substances Are Made of Atoms • Atomic Theory • Few people believed in the atomic theory ~400 B.C. • Atomic theory states that atoms are the building blocks of matter • REMEMBER: a compound is a pure substance composed of atoms of 2 or more elements that are chemically combined • Observations & experiments led to the following laws…

4. Law of Definite Proportions • States that two samples of a given compound are made of the same elements in exactly the same proportions by mass regardless of the sizes or sources of the samples • Simpler: No matter the size or source of sample of a compound; it will contain the same amount of each element it’s composed of.

5. EXAMPLE 1: Ethylene glycol (C2H6O2) (main component of automotive antifreeze) • 51.56% oxygen; 38.70% carbon & 9.74% hydrogen • Based on the Law of Definite Proportions; ethylene glycol ALWAYS contains 2-Carbon atoms, 6-Hydrogen atoms, 2-Oxygen atoms • EXAMPLE 2: Sodium chloride (NaCl) • 60.66% chlorine; 39.34% sodium

6. Brownies Example • How do we go about making brownies? • What do we rely on to make brownies?

7. Law of Multiple Proportions • States that when two elements combine to form two or more compounds, the mass of one element that combines with a given mass of the other is in the ratio of small whole numbers. • Example: CO & CO2 • Mass of CO – 28amu (C – 12amu; O – 16amu) • Mass of CO2 – 44amu (C – 12amu; 2O – 16amu) • The 2 Carbon atoms compare 1:1 where as the 2 Oxygen atoms compare 1:2.

8. Law of Multiple Proportions Cont’d

9. Law of Conservation of Mass • States that mass cannot be created or destroyed in ordinary chemical or physical changes • ALSO stated as – the mass of reactants in a reaction equals the mass of the products • Reactants  Products (Left side yields right side) • Baking Soda & Vinegar Demo

10. Atomic Theory Timeline

11. Democritus 400BC • This is the Greek philosopher Democritus who began the search for a description of matter more than 2400 years ago. • He asked: Could matter be divided into smaller and smaller pieces forever, or was there a limit to the number of times a piece of matter could be divided?

12. Atomos • His theory: Matter could not be divided into smaller and smaller pieces forever, eventually the smallest possible piece would be obtained. • This piece would be indivisible. • He named the smallest piece of matter “atomos,” meaning “not to be cut.”

13. Atomos • To Democritus, atoms were small, hard particles that were all made of the same material but were different shapes and sizes. • Atoms were infinite in number, always moving and capable of joining together.

14. This theory was ignored and forgotten for more than 2000 years!

15. Why? • The eminent philosophers of the time, Aristotle and Plato, had a more respected, (and ultimately wrong) theory. Aristotle and Plato favored the earth,fire, air and waterapproach to the nature of matter. Their ideas held sway because of their eminence as philosophers. The atomos idea was buried for approximately 2000 years.

16. Dalton’s Atomic Theory • John Dalton; 1808; an English school teacher; used the Greek concept as well as the Law of Conservation of Mass; Law of Definite Proportion; & Law of Multiple Proportion to develop his Atomic Theory. • According to Dalton; elements are composed of only one kind of atom and compounds are made from two or more kinds of atoms.

17. Also reasoned that ONLY whole numbers of atoms can combine to form compounds. • Revised the early Greek idea of atoms into scientific theory that could be tested by experiements

18. Dalton’s Atomic Theory 1. All matter is composed of extremely small particles called atoms, which cannot be subdivided, created or destroyed. 2. Atoms of a given element are identical in their physical and chemical properties. 3. Atoms of different elements differ in their physical and chemical properties. 4. Atoms of different elements combine in simple, whole-number ratios to form compounds 5. In chemical reactions, atoms are combined, separated, or rearranged but never created, destroyed, or changed

19. Data gathered since Dalton’s time has proven that not all of his theories are correct. • #1 & 2 are not always true. • Atoms can be divided into smaller particles. • Technology has allowed scientists to create and destroy atoms. • Dalton’s Theory has since been modified & expanded as scientists learn more about atoms.

20. Section 2: Structure of Atoms • Subatomic particles: electrons, neutrons, and protons

21. Thomson’s Plum Pudding Model • In 1897, the English scientist J.J. Thomson provided the first hint that an atom is made of even smaller particles.

22. Thomson Model • He proposed a model of the atom that is sometimes called the “PlumPudding” model. • Atoms were made from a positively chargedsubstance with negatively charged electrons scattered about, like raisins in a pudding.

23. Thomson Model • Thomson studied the passage of an electric current through a gas. • As the current passed through the gas, it gave off rays of negatively charged particles.

24. Thomson Model • This surprised Thomson, because the atoms of the gas were uncharged. Where had the negative charges come from? Where did they come from?

25. Thomson’s Conclusions • Thomson concluded that the negative charges came from within the atom. • A particle smaller than an atom had to exist. • The atom was divisible! • Thomson called the negatively charged “corpuscles,” today known as electrons. • Since the gas was known to be neutral, having no charge, he reasoned that there must be positively charged particles in the atom. • But he could never find them.

26. Rutherford’s Gold Foil Experiment • In 1908, the English physicist Ernest Rutherford was hard at work on an experiment that seemed to have little to do with unraveling the mysteries of the atomic structure.

27. Rutherford’s experiment Involved firing a stream of tiny positively charged particles at a thin sheet of gold foil (2000 atoms thick)

28. Most of the positively charged “bullets” passed right through the gold atoms in the sheet of gold foil without changing course at all. • Some of the positively charged “bullets,” however, did bounce away from the gold sheet as if they had hit something solid. He knew that positive charges repel positive charges.

30. This could only mean that the gold atoms in the sheet were mostly open space. Atoms were not a pudding filled with a positively charged material. • Rutherford concluded that an atom had a small, dense, positively charged center that repelled his positively charged “bullets.” • He called the center of the atom the “nucleus” • The nucleus is tiny compared to the atom as a whole.

31. Rutherford • Rutherford reasoned that all of an atom’s positively charged particles were contained in the nucleus. The negatively charged particles were scattered outside the nucleus around the atom’s edge.

32. Atomic Number & Mass Number • Atomic Number = number of protons in the nucleus of an atom; the atomic number is the same for all atoms of an element; ALWAYS whole numbers • The atomic number also tells us how many electrons there are in a neutral atom

33. Mass Number = the sum of the numbers of protons and neutrons of the nucleus of an atom • #of Neutrons = Atomic Mass - #of Protons • Elements can have varying numbers of neutrons and electrons but NOT protons.

34. Symbols for elements can be as follows: • 1/1H • 4/2He • 7/3Li • 1,4, &7 represent the mass number • 1,2,&3 represent the atomic number

35. Isotopes: an atom that has the same number of protons as other atoms of the same element BUT has different number of neutrons • Two ways of designating isotopes • H-1; H-2; H-3 • He-3; He-4 • 3/2He; 4/2He (Mass number on top/Atomic Number on bottom)

36. HOMEWORK: Worksheet; Page 89 4-6

37. Section 3: Electron Configuration • After the atomic theory was widely accepted scientists began constructing models of the atom. • After Rutherford discovered the nucleus; his model proposed electron orbitals. • Suggested that electrons revolve around the nucleus like the planets orbit around the sun.

38. Bohr Model • In 1913, the Danish scientist Niels Bohr proposed an improvement. In his model, he placed each electron in a specific energy level.

39. Bohr’s Model • Niels Bohr, a Danish physicist • Describes electrons in terms of their energy levels • According to Bohr; electrons are restricted to particular locations so far from the nucleus • An electron close to the nucleus is in the lowest energy level; the farther an electron is from the nucleus the higher the energy level the electron occupies

40. Orbitals – regions around the nucleus that correspond to specific energy levels; region where electrons are most likely to be found. • Orbitals are also sometimes called electron clouds because they do not have sharp boundaries; orbitals show where it is most likely to find an electron • Example: When a fan is running you cannot tell where any one blade is at one time; but you know it is there

41. Electrons & Light • 1900 – Scientists knew that light could be thought of as moving waves that have given frequencies, speeds, & wavelengths • Speed of Light = 2.998e8m/s • Wavelength = the distance between 2 consecutive peaks or troughs of a wave (measured in meters) • Wavelength of light varies between 105m to 10-13m

42. This range makes up the electromagnetic spectrum • Our eyes are only sensitive to a small portion of this spectrum; from 400nm (violet light) to 700nm (red light) • ROYGBIV (Red; Orange; Yellow; Green; Blue; Indigo; Violet)

44. Light is an Electromagnetic Wave • When passing light through a prism you get the visible spectrum….ROYGBIV • The electromagnetic spectrum also includes X-rays, ultraviolet light, & infrared light, microwaves & radio waves • Each form of light is referred to light; although we cannot see them.

45. Red light has a low frequency and a long wavelength (moves slower) • Violet light has a high frequency and a short wavelength (moves faster) • Frequency & Wavelength are INVERSELY related (one gets smaller the other gets longer; one gets longer the other gets smaller)

47. When high-voltage current is passed through a tube of hydrogen gas at low pressure, a lavender-colored light is seen • The spectrum of a few colors is called a line-emission spectrum • Electrons can move from a low energy level to a high energy level by absorbing energy • Electrons at higher energy levels are unstable and can move to a lower energy level by releasing energy. This energy is released as light with a specific wavelength

48. Each move from a higher energy level to a lower energy level releases different wavelengths of light

49. If an electron is in the lowest energy state possible than it is said to be in the GROUND STATE • If an electron gains energy; it moves to an EXCITED STATE • As electrons move from the excited state to the ground state certain wavelengths are released depending on the difference between the two states

50. Quantum Numbers • Present-day model of atom, electrons are located in orbitals which are also called QUANTUM NUMBERS • Quantum Numbers are regions of high probability for finding electrons • Quantum Numbers are equal to your n values; which are whole positive integers for example: 1, 2, 3, 4, etc… • As n increases so does the distance of the orbital from the nucleus