Chapter 4

1. Chapter 4 Atomic Structure

2. Objectives • Be able to define an atom • Distinguish between a molecule and a compound

3. Atoms • Definition - the smallest particle that has the properties of an element, basic unit of matter - surprising the number of atoms is small - the number of combinations these atoms make are huge ex. color print: 3 colors can make numerous colors • 119 distinct atoms as of 1999, form elements

4. Atoms Cont. • hydrogen makes up more than 90% of the atoms in the universe • 1st direct evidence was inadvertently discovered in 1827, by Scottish botanist Robert Brown while he was studying pollen • grains were in a constant state of agitation - thought they were moving life forms, later discovered a perpetual jiggling of particles known as brownianmotion; collisions between visible particles and invisible atoms

5. Atomic Symbols • each element has its own name, accompanied by a symbol • usually two/three letters (first is always capitalized) ex. Iron: Fe • Fe represents 1 atom of iron, • 2Fe represents 2 atoms of iron etc… • can also be written as Fe2

6. Molecules • Definition - the smallest unit of a substance that exhibits all the properties characteristics of that substance • two or more atoms ex. H8, O2

7. Compound • Definition - a substance that is made from two or more simpler substances and can be broken down into those simpler substances ex. H2SO4

8. Objectives • Describe ancient Greek models of of matter • List the main points of Dalton’s atomic theory and describe his evidence for the existence of atoms • Explain how Thomson and Rutherford used data from experiments to produce their atomic models

9. Democritus Theory • Greek Philosopher in the 4th century B.C - believed that all matter consisted of extremely small particles - suggested these particles are made of invisible units called atoms - term atom is derived from a Greek word meaning “unable to divide” - believed there were different types of atoms, liquids: round,smooth solids: rough, prickly - unable to provide evidence that an atom existed, therefore many people were very skeptical

10. Dalton’s Atomic Theory • John Dalton - interested in predicting the weather SO.. studied the behavior of gases in the air, concluded that a gas consists of individual particles • Evidence - masses of elements as they combined to form compounds always produced the same ratio no matter what the size of the sample ex. carbon dioxide - 1 carbon 2 oxygen: 1:2 ratio

11. Dalton’s Atomic Theory • Theory - used a Greek concept of the atom and the 3 laws to give the atomic theory a scientificbasis

12. Dalton’s Atomic Theory Cont. • 5 Principles 1. All matter is made of indivisible and indestructible atoms 2. All 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-numbersratios to form compounds

13. Dalton’s Atomic Theory Cont. 5. Chemical reactions consist of the combination, separation, or rearrangement of atoms - theory explained most of the chemical data of the day and was readily accepted - evidence since has shown the first twoprinciples are not valid; overlooked that most atoms will combined with other of their own kind - NOT discarded only modified

14. Objectives • Explain JJ Thompson’s experiment in detail • Understand Thompson’s observations and conclusions • Draw a Thomson’s model of the atom

15. Thomson’s Experiment • J.J Thomson 1st experiment - pumped most the air out of a glass tube, placed a metal plate at each end - applied a voltage to two metal plates one became positively charged:anode other became negatively charged: cathode

16. Thomson’s Experiment Cont. 2nd experiment - placed a charged metal plate on either side of the glass tube

17. Thomson’s Experiment Cont. • Observations 1st experiment - glowingray emerged between the cathode and anode 2nd experiment - charged plates caused the beam to deflect/bend - repelled from the negative - attracted to the positive • Conclusion - beam of light (stream of charged particles) - negative

18. Thomson’s Experiment Cont. • Uses - TV screens - computermonitors - radar displays *Later became known as cathode rays due to their origin now known as an electron beam • 1st to provide evidence that atoms are made of smallerparticles • revised Dalton’s model

19. Thomson’s Model * If there is a negative charge there must also be a positive charge • Why? - believes the atom is neutral • Plum Pudding - negative particles are evenly scattered throughout an atom with a positively charged mass of matter - similar to that of chocolate chip ice cream - later proved to be incorrect

20. Objectives • Explain Rutherford’s experiment in detail • Understand Rutherford’s observations and conclusions • Draw Rutherford’s model of the atom • Compare and contrast Thomson’s and Rutherford’s models

21. Rutherford’s Theory • Ernest Rutherford - a former student of Thomson came up with a more accurate picture of the atom in 1909 -oversaw the now famous goldfoilexperiment • Gold Foil Experiment • Hypothesis - alpha particles are thousands of times more massive, hence they would not be impeded as it passed through the “atomic pudding” - beam of positively charged particles, alphaparticles from a radioactive source was directed through a sheet of very thin gold foil

22. Gold Foil Experiment

23. Gold Foil Experiment Cont. • Observations - nearly all passed through undeflected and produced spots of light - some were widelydeflected, and a few bounced straightback What massive object did they hit? atomicnucleus, an extremely dense positively charged center of the atom

24. Gold Foil Experiment Cont. Why did the others pass through then?

25. Gold Foil Experiment Cont. • Conclusion - atom is mostly emptyspace - most of its mass concentrated in the central region, atomicnucleus

26. Gold Foil Experiment Cont. - the nucleus and surrounding electrons occupy only a tinyfraction of the atomic volume - diameter of an atom is generally about 10,000 times greater than the diameter of its nucleus * If the nucleus were the size of the period at the end of this sentence, the outer edges of the atom would be located some 3.3 meters away *

27. Gold Foil Experiment Cont. • Rutherford’s Model - all of the atoms positive charge is concentrated in the nucleus, which only takes up a very small amount of the atom Can we then say we are mainly empty space?

28. Objectives • Identify three subatomic particles • Understand how subatomic particle was discovered • Compare the properties of the subatomic particles • Distinguish between atomic number and mass number • Calculate the number of protons, electrons and neutrons in an atom

29. Subatomic Particles • Subatomic particles - 3 important to chemistry - protons, neutrons, electrons

30. Protons • Definition - a positively charge subatomic particle that is found in the nucleus of an atom • About Protons - proton is nearly 2000 times more massive than the electron, but equal in charge and opposite in sign to the electron - number of protons in the nucleus is electrically balanced by an equal number of electrons ex. oxygen atom: contains 8 electronsand protons: neutral atom, no net charge

31. Electron • Definition - a negatively charged subatomic particle that is found in the space outside the nucleus - name comes from the Greek word for amber - Amber: material discovered by early Greeks that was found to exhibit the effects of electrical charging ex. Ben Franklin: Key/Kite - lead others to experiment with electric currents through gases in sealed tubes

32. Neutrons • Definition - is a neutral subatomic particle that is found in the nucleus of the atom - mass almost exactly equal to that of the proton

33. Comparing Subatomic Particles

34. Atomic Number • Definition - number of protons in the atom ex. Oxygen 8p + 8n = 16 - elements are classified by this number - continues up to 119 - unique to a given element - all atoms are electrically neutral, meaning the number of electrons must equal the number of protons • this arrangement of elements by their atomic numbers makes up the periodictable • Usually located at the upperleft hand corner

35. Mass Number • Definition - the total number of protons and neutrons in the nucleus of an atom - mass number – atomic number = neutrons ex. N: mass number of 14 atomic number of 7 7 neutrons

36. Mass Number Cont. • usually found at the bottom of the atomic symbol, and sometimes found written at the bottom left of an atomic symbol ex. 16O

37. Mass Number - although a given type of atom will usually contain a certain number of neutrons in the nucleus, a small percentage will not ex. most hydrogen atoms contain no neutrons - a small percentage contain one neutron and a smaller percentage two neutrons • What do we call atoms with a different number of neutrons? - isotopes

38. Isotopes • Definition - the number of neutrons in the nucleus of a given element may vary, protons remain the same ex. H contains 1 protons (1H) H contains 1 protons and 1 neutrons (2H) deuterium H contains 1 protons and 2 neutrons (3H) tritium ex. 14C: Carbon-14

39. Atomic Mass • Definition - mass of an atom in atomic mass units (amu) - atoms have very little mass - equal to 1/12th of the mass of carbon - often an average mass - weighted mass • AMU or the Dalton (Da) - equal to 1.6605402 x 10-27 kg

40. Atomic Mass Number Cont. ex. 99% of all carbon atoms are the isotope containing 6 neutrons, the remaining 1% is the heavier isotope containing 7 neutrons, which raises the average mass of carbon from 12.000 to 12.011

41. Objectives • Describe Bohr’s model of the atom and the evidence for energy levels • Explain how the electron cloud model represents the behavior and locations of electrons in atoms

42. Bohr’s Model of the Atom • Bohr’s Model - an early conceptualmodel of the atom - classic planetary model in which electrons whirl around the small but dense nucleus: like planets orbiting the Sun - developed by the Danish physicist Niels Bohr in 1911 - each electron has a certain energy that is determined by it’s path around the nucleus - explains how atoms (lose) emit or (gain) absorb energy resulting energy: energy level

43. Bohr’s Model: Energy Levels • Definition - any of the possible energies an electron may have in an atom • Evidence - measured amount of energy gained or lost - when energy is lost we often see it as a light ex. fireworks - 1925 this model no longer explained all observations being made by scientists and new models were created ex. Electron cloudmodel

44. Electron Cloud Model • visual model of the most likely locations for electrons in an atom • Orbitals - a region in an atom where there is a high probability of finding an electron ex. propeller on a helicopter (you know its there you see a blur, can’t pinpoint exact location) - 4 orbitals s – 2 houses 2 e- p – 3 houses 6 e- d – 5 houses 10 e- f – 7 houses 14 e-

45. Orbitals - electrons occupy the lowest energy levels first - electrons in the outermost energy levels of an atoms are called valenceelectron

46. Electron Configurations • Definition - arrangement of electrons in the orbitals of an atom (similar to seating assignments on an airplane) - when all electrons at their lowest energies this is called groundstate