ATOMIC STRUCTURE Name: Period: Date: I. LAW vs. THEORY: 1) = a generalization of scientific observations that what happens (does explain) 2) (model) = a set of assumptions used to explain observations and predict new observations a) Can be truly proven 100% correct. b) Inevitably change (and must sometimes be abandoned) as more information becomes available. c) Considered successful when they observations and, more importantly, new observations. II. THE ATOM: FROM PHILOSOPHICAL IDEA TO SCIENTIFIC THEORY: 1) EARLY ATOMIC THEORY: a) 400 B.C- (Greek philosopher) Coined the term (meaning indivisible) to describe the smallest particles of matter Did not experiment 2) JOHN DALTON (EARLY 1800 S): a) Father of the b) Dalton s Atomic Theory helped explain measurable observations and successfully predict new observations (while stimulating additional research from Dalton and other scientists). c) Assumptions of Dalton s Atomic Theory (proposed by Dalton in 1803): i. All matter is composed of 1
ii. All atoms of the same element are in size, mass, other properties. Atoms of different elements differ in size, mass, and other properties. iii. Atoms cannot be subdivided, created, or destroyed iv. : Atoms of different elements combine in simple whole-number ratios to form chemical compounds. v. In, atoms are combined, separated, or rearranged. d) Dalton used his Atomic Theory to help correctly new observations leading to the Law of Multiple Proportions. e) = The masses of one element that combine with a constant mass of another element to form more than one compound are in the ratio of small whole numbers Carbon monoxide and carbon dioxide contain oxygen in a 1:2 ratio. of carbon reacts with of oxygen to form carbon monoxide, CO. of carbon will react with of oxygen to form carbon dioxide, CO 2. f) Dalton s Atomic Theory explained many observations and correctly predicted many additional observations. It correctly predict all new observations (no theory ever does). g) Dalton s Atomic Theory has been to explain the new observations: Atoms of one element can have masses ( ). Atoms subdivided (but not in a chemical reaction or physical change) Nuclear reactions An atom of one element changed into an atom of another element (but not in a chemical reaction or physical change) Nuclear reactions h) As with all theories, Atomic Theory has been changed and expanded over time (and will continue to change and expand) to explain new observations. 2
3) CATHODE-RAY TUBE EXPERIMENTS (late 1800 s): = A glass tube containing a gas at a very low pressure which contains a negative electrode (cathode) and a positive electrode (anode). When high voltage is applied, a cathode ray passes from the cathode to the anode causing the low pressure gas in the tube to glow (different gases glow different colors). a) Early cathode-ray tube experiments: Proposed that a cathode ray consists of tiny particles with mass and they are negatively charged. b) discovered the Used early cathode-ray tube experiments and some of his own findings to support the hypothesis that electrons are negatively charged particles Used a - as voltage across the tube was increased, a beam of light (cathode ray) became visible = beam of electrons seen b/c of excited gas Found that the beam was by both magnetic and electrical fields OBSERVATION: Noticed that the cathode rays were attracted to the positive electrode, called the anode. What conclusion did Thompson have? The cathode rays were made up a very small charged particle OBSERVATION: Measured the bending of the path of the cathode rays and was able to determining the ratio of an electron s charge to its mass. Found that the ratio was always the, regardless of the metal used. 3
What conclusion did Thompson have? Concluded that the negatively charged particles were much than the lightest know atom (hydrogen), which meant that atoms had a! c) Later were also discovered using a modified cathode-ray tube d) Thomson s Model of the Atom ( ) Thomson postulated that an atom consisted of a diffuse cloud of positive charge with the negative embedded randomly in it. This represented a major from Dalton s model of atoms as indivisible. The existence of the electron raised new questions: if electrons are part of all matter and they possess a charge, how can all matter be? Also, if the mass of an electron is so small, what accounts for the rest of the in a typical atom? o More experiments to come 4) ERNEST RUTHERFORD & THE GOLD FOIL EXPERIMENT a) Rutherford discovered the using alpha particles & gold foil = a relatively large positively charged particle b) Gold-Foil Experiment (around 1908-1909): Thin was bombarded by alpha particles and the path of the alpha particles was charted after they passed through the gold foil. c) Expected results: The massive (positively charged) were expected to crash through the gold foil with little or deflections; there was nothing in Thomson s model of the atom to cause anything more than minor deflections of the alpha particles. 4
(a) When a beam of alpha particles is directed at a thin gold foil, most particles pass through the foil undeflected, but a small number are deflected at large angles and a few bounce back toward the particle source. (b) A closeup view shows how most of an atom is empty space and only the alpha particles that strike a nucleus are deflected. d) OBSERVATION: Some alpha particles were deflected at large angles, and some were redirected backward (!) What conclusion did Rutherford draw from this evidence? The POSITIVE particles of the atom must, but instead must be concentrated at the center of the atom-the ( ). e) OBSERVATION: Most of the alpha particles passed through the gold foil with few deflections. What conclusion did Rutherford draw from this evidence? Most of the alpha particles did not hit anything and passed straight through the gold atoms so therefore, most of the volume of an atom consists of. f) Rutherford s Nuclear Model Explained the nature of matter: the positive charge of the nucleus balances the negative charge of the electrons. Suggested that the electrons travel the positively charged nucleus. The early nuclear model did not account for all the atom s By 1920, Rutherford refined the concept of the nucleus and concluded that the nucleus contained positively charged particles called. 5
5) JAMES CHADWICK (1932) a) Discovered the III. THE MODERN MODEL OF THE ATOM: 1) WHAT IS THE DIFFERENCE BETWEEN AN ATOM AND AN ELEMENT? a). = a substance that cannot be broken down to other substances by a chemical reaction. b). : the smallest particle of an element that can exist either alone or in combination with other atoms. 2) SUBATOMIC PARTICLES OF AN ATOM a) = a particle smaller than an atom Ex: proton, neutron, electron b) An atom is composed of subatomic particles including protons, neutrons, and electrons (plus scientists have determined that protons and neutrons have their own structures and they are composed of quarks- these particles will not be covered since scientists do not yet understand if or how they affect chemical behavior). 6
c) SUMMARY OF SUBATOMIC PARTICLES Particle Symbol Electron Location In the space surrounding the Relative Electron Charge Relative Mass Actual Mass (g) 1/1840 9.110 10 28 g Proton In the 1 1.673 10 24 g Neutron In the 1 1.675 10 24 g d) give the nucleus the positive charge & the identity of an atom e) Protons and neutrons have about the mass and are over 1840 times more massive than electrons. f) Most of the mass of an atom is located in the (protons & neutrons). (Nuclear forces hold the particles of a nucleus together. These forces only act over a very short range.) g) are electrically b/c of the number of protons EQUALS the number of electrons. h) = the positively charged, dense central portion of an atoms Contains nearly all of the atom s, but takes up a small fraction of its 3) ATOMIC NUMBER & AVERAGE ATOMIC MASS 19 K Potassium 39.098 : the number of protons in the nuclei Also represents the number of that an atom has since atoms are electrically neutral Ex: Atomic number of Na is 11; Na has 11 and 11 7
: the weighted average of the atomic masses of naturally occurring isotopes of an element. = atoms of the same element with different number of NEUTRONS, therefore, different masses. lists average atomic mass Weighted averages account for the percentages of each isotope of a given element Important b/c they indicate relative mass relationships in chemical reactions Ex: How average atomic mass is calculated 4) ISOTOPES = atoms of the same element with different number of NEUTRONS, therefore, different masses. Isotopes of an element all have the SAME NUMBER of and Named by their = the total number of protons and neutrons in the nucleus of an isotope. Mass Number = # Protons + # Neutrons Ex: H-1 H-2 H-3 8
Example 1: ALL carbon atoms have how many protons? Most carbon atoms have 6 neutrons. What is their mass number? Some carbon atoms have 8 neutrons. What is their mass number? C-12 and C-14 are isotopes of carbon o Both isotopes have electrons & protons, but differ in the # of!! Example 2: How many neutrons are in a sodium-23 atom? Example 3: How many protons, electrons and neutrons are there in an atom of chlorine-37? Hyphen Notation: Write the hyphen notation for a hydrogen isotope with a mass number of 3. or Nuclear Symbol Notation: Write the nuclear symbol for a hydrogen isotope with a mass number of 3. Complete the following table. Remember: o atomic number = number of protons = number of electrons o mass number = number of protons + number of neutrons Nuclear Symbol Atomic Number Mass Number Number of Protons Number of Neutrons Number of Electrons 30 70 108 74 87 50 9
= SI unit for amount of substance A mole simply represents a, much in the same way that a dozen represents a set of twelve. Ex: Just as you can have a dozen cans of soda or a dozen donuts, you can have a mole of stars or a mole of water molecules. So a dozen eggs represents 12 eggs, a mole of carbon represents carbon atoms is also called = the number of particles in exactly one mole of a pure substance. There are atomic mass units (amu) in gram!! This conversion allows us to use the masses listed on the periodic table for both atomic mass and molar mass. = is the mass of one atom, is measured in atomic mass units (amu) Mass of C-12 is arbitrary assigned a mass of 12 amu 1/12 the mass of a carbon-12 atom is Mass of any atom is expressed relative to the mass of one atom of carbon-12 = is the mass of one mole of atoms or molecules, is measured in grams. Numerically equal to the atomic mass 1 mole of a substance is 1 molar mass of that substance Ex: 1 mole of nitrogen is equal to 14.007g of nitrogen! Element & Symbol Carbon (C) Helium (He) Copper (Cu) Potassium (K) Atomic & Molar Masses Atomic Mass Mass of 1 Atom 12.011 amu 4.0026 amu 63.546 amu 39.1098 amu Molar Mass Mass of 6.02 X 10 23 atoms REMEMBER: one mole of atoms contains 6.02 X 10 23 atoms. One mole of molecules contains 6.02 X 10 23 molecules. One mole of formula units contains 6.02 X 10 23 formula units. One mole of ions contains 6.02 X 10 23 ions. 10
Conversions: Using Charts! Ex: Conversion factors: 12 inches = 1 foot 3 feet = 1 yard Set up conversion charts so your units will cancel out! You multiple across the top of the chart and divide by what is on the bottom of the chart. 1) How many feet are in 60.72 yards? 2) How many yards are in 678.35 feet? 3) How many inches are in 964.8 feet? 4) How many feet are in 1,743.2 inches? Mrs. H s Rounding: We will use 6.02 X 10 23 atoms We will round the molar mass values from the Periodic Table (0.5 + round to the whole #) o Ex: Phosphours 30.974g would be rounded to 31g 11
MOLESMASS (grams) o KEY IDEA: It is important to know the following conversion: The of an element is equal to the mass of of atoms of that element. Example: 1 mole of zinc is equal to 65.39 grams of zinc! Use the factor label method (units will cancel out!) Ex: What is the mass in grams of 3.5 mol of the element copper (Cu)? 3.50 mol Cu 64g Cu 1 mol Cu = 224g Cu 1. What is the mass in grams of 2.25 mol of iron (Fe)? 2. What is the mass in grams of 0.375 mol of K? 12
MASS (grams) MOLES 1 mol of an element = the molar mass of that element Ex: A chemist produced 11.9g of Al. How many moles of Al has been produced? 11.9 g Al 1 mol Al 27g Al = 0.441 mol Al 1. How many moles of Ca are contained in 5.0g of Ca? 2. How many moles of Au are contained in 3.6 x 10-10 g of Au? ATOMS MOLES of an element = of that element Ex: How many moles of Ag are in 3.01 x 10 23 atoms? 3.01 x 10 23 Ag atoms 1 mol Ag 6.02 x 10 23 Ag atoms = 0.5 mol Ag 1. How many moles of Pb are equivalent to 1.5 x 10 12 atoms? 2. How many moles of tin are equivalent to 2500 atoms? 3. How many atoms of Al are contained in 2.75 mol? 13
The Mole, Molar Mass, # of Atoms Practice Problems What is the conversion factor between moles and atoms? What is the conversion factor between moles and mass? 1. What is the mass in grams of 4.25 mol of lithium? 2. How many moles of Mg are equivalent to 3.011 x 10 23 atoms? 3. How many moles of Ca are contained in 150g of Ca? 4. How many atoms of K are equivalent to 5.83 mol? 5. How many moles of Ca are contained in 40.1g of Ca? 6. How many moles of Pb are contained in 3.25 x 10 5 g of Pb? 7. How many moles of Zn are equivalent to 2.25 x 10 25 atoms? 14
8. What is the mass in grams of 6.52 mol of the C? 9. How many moles of Fe are contained in 2.65g of Fe? 10. How many atoms of Na are equivalent to 1.50 mol? 11. What is the mass in grams of 1.38 mol of the N? 12. How many moles of O are contained in 4.50 x 10-12 g of O? 13. What is the mass in grams of 8.075 mol of the Au? 15
History of the Atom Worksheet Name: Period: 1) What is the difference between a scientific law and theory? 2) Who was the first person to come up with the idea of an atom? 3) Who is known as the Father of the Modern atomic theory? 4) List 2 of the 5 statements that sum up Dalton s atomic theory. 5) Is modern atomic theory the same as Dalton s atomic theory? 6) What 3 modifications to Dalton s atomic theory have been necessary? 7) Who discovered electrons? What did he use to discover electrons? 8) a) How did cathode rays behave near positively charged objects? b) What did these observations suggest about cathode rays? 9) What are cathode rays? 10) What did Ernest Rutherford discover? 16
11) During the gold foil experiment some alpha particles (positively charged) passed through the gold foil. What did this suggest? 12) During the gold foil experiment some alpha particles (positively charged) bounced back toward the main source. What did this suggest? 13) What is the name of Thomson s model of the atom? 14) Describe Thomson s model of the atom? 15) Rutherford s model of the atom is called the Nuclear Model. Describe how this model is different from Thomson s. 17
Worksheet: Parts of An Atom 1) What does the atomic number represent? 2) What does the mass number represent? 3) Complete the following table. Also, properly label each nuclear symbol with the correct atomic number and mass number. Nuclear Symbol Atomic Number Mass Number Number of Protons Number of Neutrons Number of Electrons 19 39 78 80 34 16 10 5 137 81 40 18 5 4 9 10 56 26 37 17 30 35 34 29 18
4) What is the difference between an atom and an element? 5) Compare and contrast the three types of subatomic particles in terms of location in the atom, mass, and relative charge. 6) Atoms are electrically neutral. What does this mean in terms of the number of protons and electrons in an atom? 7) The following are isotopes of oxygen: O-17 & O-18. In the table below, fill in how many electrons, protons and neutrons each isotope has. Isotope # ELECTRONS # PROTONS # NEUTRONS O-17 O-18 8) Write the nuclear symbol AND the hyphen notation for each isotope described below: a) atomic number = 2 and mass number = 4 b) atomic number = 4 and mass number = 9 c) atomic number = 30 and mass number = 65 19
Atomic Structure: Unit Review Worksheet Name: Period: Date: 1. was the first person to come up with the idea of an atom. 2. is known as the Father of the Modern Atomic Theory. 3. discovered the neutron. 4. discovered electrons. 5. discovered the nucleus. 6. are a beam of electrons 7. Cathode rays have a charge 8. Describe how electrons were discovered? 9. Describe how the nucleus was discovered? 10. Experiments & Findings: a) Observation: During the cathode ray experiment, cathode rays were attracted to the positive charge of a magnet. Finding: What did this suggest? 20
b) Observation: During the gold foil experiment some alpha particles (positively charged) passed through the gold foil. Finding: What did this suggest? c) During the gold foil experiment some alpha particles (positively charged) bounced back toward the main radiation source. Finding: What did this suggest? 11. Compare and contrast the three types of subatomic particles in terms of location in the atom and relative charge. a) - found in the nucleus; positive charge b) found in the nucleus; neutral charge c) found in the electron clouds; negative charge 12. How are atoms electrically neutral? 13. Atoms of the same element with different number of NEUTRONS are called 14. What are isotopes named by? 15. What is the hyphen notation for the following: atomic number = 6; mass number = 14 16. What is the nuclear symbol notation for the following: atomic number = 26; mass number = 56 17. The sum of protons and neutrons in an atom equals the 18. The number 34 in the name sulfur-34 represents 19. The weighted average of the atomic masses of naturally occurring isotopes of an element is the 20. The si unit for an amount of a substance is the 21. is the mass of one mole of atoms or molecules, is measured in grams 22. 1/12 the mass of a carbon-12 atom is 23. What is the conversion factor between moles and grams? 21
24. What is the conversion factor between moles and atoms? 25. What is the mass in grams of 4.89 mol of the O? 26. How many moles of Au are equivalent to 9.62 x 10 18 atoms? 27. How many moles of B are contained in 236.9g of B? 28. How many atoms of Rb are equivalent to 8.51 mol? 29. Complete the following table. Also, properly label each nuclear symbol with the correct atomic number and mass number. Nuclear Symbol Atomic Number Mass Number Number of Protons Number of Neutrons Number of Electrons 35 80 15 16 137 56 30. What is the difference between a scientific law and a theory? 22
Unit Learning Map (6 days): Atomic Structure Mrs. Hostetter Class: Academic Chemistry A - Grade 11 Unit Essential Question(s): How do scientists view the atom? Optional Instructional Tools: Guided Notes Lab Materials: Al Atom Concept Concept Concept Concept Theory Model of an atom The mole Lesson Essential Questions: Lesson Essential Questions: Lesson Essential Questions: Lesson Essential Questions: How did scientists contribute to the modern atomic theory? How has the model of the atom evolved? What is a mole? Vocabulary: Vocabulary: Vocabulary: Vocabulary: Scientific Law Theory Democritus John Dalton J.J. Thompson Ernst Rutherford James Chadwick Atom Element Law of Multiple Proportions Cathode ray Nucleus Protons Neutrons Electrons Isotopes Atomic number Average atomic mass Mass number Nuclear symbol notation Hyphen notation Mole Molar mass 23
Atomic Structure Vocabulary: 3) Scientific law- a generalization of scientific observations that describes what happens (does not explain) 4) Theory - a set of assumptions used to explain observations and predict new observations 5) Democritus- Greek philosopher that coined the term atom (meaning invisible) to describe the smallest particles of matter 6) John Dalton- Father of the Modern Atomic Theory -All matter is composed of atoms -All atoms of the same element are identical in size, mass and other properties. -Atoms cannot be subdivided, created or destroyed -Law of Multiple Proportions -In chemical reactions, atoms are combined, separated or destroyed 7) J.J. Thomson- discovered the electron using a cathode ray tube 8) Ernst Rutherford- discovered the nucleus through the gold foil experiment 9) James Chadwick- discovered the neutron 10) Atom- the smallest particle of an element that can exist either alone or in combination with other atoms 11) Element- a substance that cannot be broken down to other substances by a chemical reaction 12) Law of Multiple Proportions- Atoms of different elements combine in simple whole-number ratios to form chemical compounds 13) Cathode ray- beam of electrons seen because of excited gas 14) Nucleus- positively charged, dense central portion on an atom 15) Protons- gives the nucleus a positive charge, determines the identity of the atom, and positive charge is equal to the negative charge of the e- 16) Electrons (e-)- move nearly at the speed of light and form an electron cloud around the nucleus 17) Isotopes- atoms of the same element with different number of neutrons, therefore, different masses. 18) Atomic number- number of protons (and electrons) in the nucleus of an atom 19) Average atomic mass- the weighted average of the atomic masses of naturally occurring isotopes of an element 20) Mass number- total number of protons and neutrons in the nucleus of an isotope 21) Hyphen notation element symbol followed by the mass number; Ex: He-4 22) Nuclear symbol notation- mass number is on top; atomic number is on the bottom; Ex: 23) Mole- SI unit for amount of substance 24) Molar Mass- mass of one mole of atoms or molecules measured in grams 24