Chapter 4 Atomic Structure

1. Chapter 4Atomic Structure Ms. Riggins Lawndale High School

2. Section 4.1 - Defining the Atom • The discussion of atoms began as a philosophical concept in ancient Greece and India over 2000 years ago The discussion entered the scientific community in the early 1800s

3. Democritus’ Philosophy • The Greek philosopher Democritus was the first to suggest the existence of atoms (400BC) • He theorized that everything is composed of “atoms” • He said that atoms were indivisible, always in motion, and had a lot of empty space between them

4. Democritus’ Philosophy He also said that there are an infinite number and kind of atoms with different shapes and sizes He believed that iron atoms were solid and strong with hooks He believed that water atoms were smooth and slippery He believed that salt atoms were sharp and pointed because of their taste

5. Democritus Connecting Atoms • About 2000 year later, John Dalton used experiments to transform Democritus’ ideas on atoms into a scientific theory He believed that air atoms were light and whirly He also believed that atoms had hooks to hold them together

6. Defining the Atom • We now know that ALL matter is composed of particles called atoms (can’t see it, but it’s there) • Atoms – the smallest particle of an element that retains it identity in a chemical reaction

7. John Dalton • In 1766, Dalton ran a series of experiments with many different types of elements to formulate his atomic theory

8. Dalton’s Atomic Theory (1.) All elements are composed of tiny indivisible particles called atoms. (from Democritus) (2.) Atoms of the same element are identical. The atoms of any one element are different from those of any other element. (3.) Atoms of different elements can physically mix together or can chemically combine to form compounds.

9. Dalton’s Atomic Theory (4.) Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.

10. Section 4.2 – The Structure of the Atom • There are three kinds of subatomic particles in an atom: electrons, protons, and neutrons In 1897, J. J. Thomson discovered the electron (negatively charged subatomic particles) by using a cathode ray and came up with the Plum Pudding Model

11. Sir Joseph John Thomson • Until the late 1900s, scientists believed that atoms were indivisible J.J. Thomson proved them wrong when he discovered the electron He also discovered the concept of isotopes (we’ll talk about this in the next section)

12. Plum Pudding Model Thomson’s plum pudding model stated that electrons were embedded in a positive charge (like fruit in a Jello cup)

13. Concept of the Proton Atoms were known to be electrically neutral, which meant that there had to be some positively charged matter to balance the negative charges

14. Millikan’s Oil Drop Experiment • In 1909, Robert Millikan did an experiment to find the charge and mass of an electron by comparing the gravity force versus the buoyant force on a tiny drop of oil suspended between two metal electrodes • In 1916, he reported that the mass of an electron is 1/1840 the mass of a hydrogen atom

15. Ernest Rutherford (student of Thomson) • In 1911, Rutherford proposed that atoms had a positive charge concentrated in the nucleus Rutherford’s model became known as the planetary model of atoms In 1917, he was accredited for the first person to split the atom

16. Rutherford’s Gold-Foil Experiment Ernest Rutherford’s experiment disproved the plum pudding model of the atom and suggested that there was a positively charged nucleus (central core of an atom)

17. Conclusion of Rutherford’s Experiment • Atoms are mostly empty space, thus explaining the lack of deflection of most of the alpha particles • All the positive charge and almost all the mass of an atom are concentrated in a small region (nucleus) • Nucleus – tiny central core of an atom composed of protons and neutrons • Electrons are distributed around the nucleus and occupy almost all the volume of the atom (marble and football stadium)

18. Structure Of An Atom So by this point, we know that protons and neutrons are located in the nucleus and electrons are around the outside of the nucleus

19. Properties of Subatomic Particles

20. Summary of Atomic Theory

21. Section 4.3 – Distinguishing Among Atoms • Elements are different because they contain different numbers of protons. Atomic Number - the number of protons in the nucleus of an atom *Remember since atoms are electrically neutral, the number of protons equals the number of electrons (until we get to Chp. 5)

22. Quick Practice… How many protons and electrons are in each atom? 1. Fluorine 2. Calcium 3. Aluminum 4. Boron 5. Neon 6. Magnesium

23. Mass Number Mass Number – the total number of protons PLUS neutrons in an atom Therefore…the number of neutrons in an atom is the difference between the mass number and the atomic number # of Neutrons = Mass # – Atomic #

24. Atomic Mass Units • Atomic Mass Unit (amu) is 1/12 the mass of a carbon-12 atom Atomic mass is the weighted average of the atom in naturally occurring samples of the element

25. Shorthand Notation(You need to know this notation) Mass Number Au 179 79 Atomic Symbol Atomic Number

26. Practice Shorthand Notation… How many protons, neutrons, and electrons are in each atom? Atomic # Mass # • Beryllium (Be) 4 9 • Neon (Ne) 10 20 • Sodium (Na) 11 23 How many neutrons are in each atom? 1. Carbon-12 2. Fluorine-19 3. Sulfur-32 6 10 16

27. Isotopes Isotopes – atoms that have the same number of protons, but different numbers of neutrons (also different mass numbers) Write the following isotopes of oxygen: 1. Oxygen-16 2. Oxygen-17 3. Oxygen-18

28. Atomic Mass Atomic Mass – weighted average mass of the atoms in a naturally occurring sample of the element In order to calculate the atomic mass of an element: (1.) Multiply the mass of each isotope by its natural abundance (2.) Add the products together

29. Let’s practice… Calculate the atomic mass of the following element, X The isotope 10X has a mass of 10.012amu and a relative abundance of 19.91%. The isotope 11X has a mass of 11.009amu and a relative abundance of 80.09%. ANSWER = 10.810amu

30. More Practice… • The element copper has naturally occurring isotopes with mass numbers of 63 and 65. The relative abundance and atomic masses are 69.2% for mass = 62.93amu, and 30.8% for mass = 64.93amu. Calculate the average atomic mass of copper. • 2. Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative abundance of 78.92amu (50.69%) and 80.92amu (49.31%).

31. Preview of the Periodic Table Periodic Table – an arrangement of elements in which the elements are separated into groups based on a set of properties Period – horizontal rows of the periodic table (there are 7) Group/Family – vertical columns of the periodic table • Elements within a group have similar chemical and physical properties

32. Homework Chapter 4 Assessment Page 122 #’s 34 – 55, 59, 61, 64, 65, 71, 78, 81, 85, 88