Chemistry I Periodic Properties and Periodicity Class Notes
|
|
- Madlyn Francis
- 5 years ago
- Views:
Transcription
1 Chemistry I Periodic Properties and Periodicity Class Notes History - Periodic Table Periodic Properties Review of electron configuration & quantum numbers I. Know what valence electrons are -those electrons involved in chemical reactions - most commonly are the outermost electrons, but may involve more (like transition metals) II. Degenerate - term for orbitals within a subshell that have the same amount of energy. ex. Px Py Pz all have the same energy -an electron may enter any one of them, or move from one to another with no energy gain or loss. III. Electron cloud - shape is based on the probability distribution curve for the location of the electron IV. In addition to number designations ( n = 1, 2, 3,...), shells can also be designated by the letters: K, L, M, N,... V. Diamagnetic vs. Paramagnetic - arrangement of electrons within an atom - whether there all electrons are paired, or if there are unpaired electrons. diamagnetic - no unpaired electrons (note: this does not simply mean that there are an even number of electrons!) paramagnetic - at least one unpaired electron VI. Isoelectronic - same number of electrons in species ex. the chloride ion, the sulfide ion, and the potassium cation all have the same number of electrons a neutral argon atom. VII. Cosmic Rays - higher in energy than even gamma rays - source - stellar explosions? ============================================================================= History of the Periodic Table Development of the Periodic Law 1770's - discovery of oxygen by Lavoisier start of modern chemistry - within the next 60 years over 50 elements are separated and identified as being elemental substances. (by 1830, 56 elements had been identified) Chemists started looking for grouping methods first separation was metals from nonmetals -still common today Dobereiner - grouped elements in triads -Law of Triads - elements can be grouped so that the "middle element" has the "average properties" of the extremes. ex. Si (MW =28) and Sn (MW = 119) average MW = 73.5 Molecular Weight of Ge is only rough approximations
2 s A. Newland - Law of Octaves - showed recurring similarities of physical properties for each eighth element - recurring based on increasing atomic weight - elements showed a similarity to those above and below it ** Noble gases had not been discovered yet, so the eighth element after Na was K ** B. de Chancourtois - produced a spiral arrangement - first attempt at putting all elements in one arrangement 1869 Meyer (German) and Mendeleev (Russian) - both worked independently and came up with similar idea at the same time - within months of each other -question of publication Mendeleev - father of periodic table 1. mainly because his table was based on both physical and chemical properties; Meyer's table was based on physical similarities 2. took bold steps to predict that "gaps" should be left in table to allow for elements that had not yet been discovered. 3. predicted some physical and chemical properties of elements as yet undiscovered. One notable example: Germanium - he called it eka-silicon (after silicon) predicted (1871) disc. by Winkler (1886) Property Mendeleev Prediction Actual appearance gray grayish-white atomic weight density specific heat Formula of oxide EO2 GeO2 Mendeleev made predictions for about several elements - books disagree - some say as few as 4, while others say as many as 14. Remember that the noble gases had not been discovered yet. Sir William Ramsey (Cambridge) discovered all inert gases from No longer called inert gases because of the synthesis of compounds with the heavier elements. These were "stuck" on the end of the table - mainly as a place to put them. Note that the date is only around the time of the discovery of the electron - before nuclear composition was known. *** Periodic Table was based on increasing atomic weights. ** The Modern Periodic Table is based on increasing atomic number - progressive increase in the number of protons. This was achieved by the work of Moseley and X ray diffraction
3 -3- Periodic Properties: Work of G.N. Lewis - proposed the idea of the octet rule even before the electron configuration theory was proposed by Bohr. Stability of an element was related to its electronic stability Effective nuclear charge and shielding discuss (see additional reading on pg. 10) Nuclear charge is the charge based on the number of protons in the nucleus. Because the atom has shells of electrons, some of these electrons shield the outer electrons of the nuclear charge. The actual (or effective) charge on valence electrons is called the effective charge. Zeff = Z - S S = # of inner shell electrons Z = atomic number = # of protons Note that the effective charge on magnesium is greater than the effective charge on sodium Zeff for Na = 1 Zeff for Mg = 2 Sodium has 1 proton pulling on 1 electron While the ratio of electron: proton is 1:1 in both sodium and magnesium, in each magnesium atom each electron feels the positive pull of two electrons!... Divisions within the periodic table: Groups or families - vertical columns - similar in chemical properties Series or Periods - horizontal rows - show progressive variation 1. I-A alkali metal - very active metals oxid.# = +1 "alkali" from ashes soda ash = sodium carbonate potash = potassium carbonate 2. II-A alkaline earth metals - active metals - found in earthen or mineral form. 3. VI-A - chalcogens 4. VII-A - halogens or halides 5. VIII or 0 inert gases, noble gases 6. I-B - coinage metals - because of unreactivity 7. VIII-B - iron family Note: except for Fe, Co, and Ni, the bottom two rows of this group (metals #44, 45, 46, and #76, 77, 78) are also known as the noble metals. They are called this because of their lack of reactivity. It is interesting to note that gold is not included in this group, because it is probably one of the least reactive of all metals! 8. A elements - called representative elements 9. B elements - called transition metals those elements with "d" subshells being filled are transition those elements with "f" subshells being filled are "inner" transition metals Inner transition metals - lanthanide series (sometimes called rare-earth) -actinide series - those past U (# 93 and beyond) called transuranium elements Some similarity between I-A and I-B true for all A and B groups. Shielding effect minimizes periodic variations of properties for the transition and inner transition metals since electrons are being put in inner subshells and outer "valence" electrons are not changed. *** NOTE: Hydrogen could actually be grouped in either I-A or VII-A
4 -4- Periodic Properties: - atomic radius/ volume - ionic size compared to atomic size - ionization energy - electron affinity - electronegativity - density - colors - melting points/ boiling points - metallic character - properties of metals, nonmetals, metalloids I. Atomic/Ionic Radius/Volume - atomic radius - measured from the nucleus to the outermost electron - not necessarily the highest energy electron. - volume is directly related to the radius - atom is considered to be spherical Periodic trends: - as we move down the table within a group, the atomic radius increases because we are filling shells that are successively farther away from the nucleus. This increase is most noticeable for the first four shells, but drops off as we move even further out. - as we move across the periodic table within a period, the radius tends to decrease. This decrease is due to the increased effective nuclear charge. Within a period, the distance to the outer shell is not increased, but the number of protons in the nucleus is increasing. - 'lanthanide contraction" -note that the radii of the 5d elements are similar to those of the 4d elements - they don't show the expected increase in size - this is due to the increased Zeffective of the element Ionic Radius For cations, the ions are smaller than the atoms. ex. Na + is smaller than the neutral Na atom For anions, the ions are larger than the atoms. ex. Cl - is larger than the neutral Cl atom Reasons: Cations are smaller than the neutral atoms because (1) the cation has lost its outermost electrons, making the new outermost electrons those of the previously filled shell; (2) the cation now has an excess of protons, so all electrons are pulled closer to the nucleus. Anions are larger than the neutral atoms because (1) the anion has gained electrons in order to achieve an octet configuration, making it more stable. This filled stability increases the electron repulsion, so electrons move further away; (2) the anion now has an excess of electrons, so all electrons are able to move away from the nucleus. Be able to compare ions to their isoelectronic noble gas in terms of size.
5 -5- II. IONIZATION ENERGY (IONIZATION POTENTIAL) - amount of energy required to completely remove an electron from an atom or ion. Important points: 1. Energy must always be added to remove electrons. 2. May have as many Ionization Potentials as there are electrons in a species. ie. Na would have 11 I.E's. + - M(g) > M(g) + 1 e H is a positive value (endothermic) Note: 1. I.E. increases as you go across a period 2. I.E. decreases as you go down a family or group Lowest Ionization Energy - bottom left corner Highest Ionization Energy - top right corner Reminder: if an atom has a filled or half-filled subshell (or shell), there is a notable irregularity due to greater shielding and/or increased electron repulsion. Periodic properties will reflect these changes (jumps). There are some noticeable irregularities: 1. from II-A to III-A 2. from V-A to VI-A 3. at VIII (inert gases) 4 from II-B to III-A These are in units of kcal/mole: 1) Na (119); ** Mg (176); ** Al (138) K (100); ** Ca (141); ** Ga (138) 2) for C (260); ** for N (335) **; for O (314) Note that the magnitude of the half-filled irregularity decreases as n increases ex. for 4th shell Ge (192); As (226); Se (225) 3) F (402) Ne high at 497 (not actually an irregularity - predictably high value because of stable octet configuration) 4) Ni (176); Cu (178); ** Zn (217); ** Ga (138) NOTE also the similarities of I.E. for transition metals For 4th series, elements show an average jump of about 50 kcal - for transition metals in row #4 Sc (151); Ti (157); V (155); Cr (156); Mn (171)? reason for small jump? THE IONIZATION ENERGIES FOR THE LANTHANIDE SERIES ARE EVEN CLOSER
6 -6- What we have been discussing are the first ionization energies; there is an ionization energy for each electron in an atom. The electron with the highest ionization energy is the one closest to the nucleus 2 reasons 1. those closest to the nucleus have the farthest distance to go to be separated from an atom. 2. as successive electrons are removed, with the number of protons staying the same, the electrons must overcome progressively larger nuclear charges. Comparison of different ionization energies: all measurements are in kcal/mol 1st 2nd 3rd 4th Na Mg Al ==========================================================================. III. Electron Affinity Electron affinity: Pay close attention because my definition and values are different from those in the book. Electron affinity: A measure of the "affinity" that an atom has for an electron. The number value is an indication of the atom's preference for the electron. In contrast to ionization energy, the values of electron affinity may be positive, negative, or zero. Even the magnitude of the positive values indicates how strongly the atom wants to accept an electron PERIODIC TRENDS: 1. Metals have a tendency to lose electrons; nonmetals tend to gain electrons. Nonmetals would be expected to have a high electron affinity. 2. As we move across the periodic table, we move from metals to nonmetals. The electron affinity should increase as we move from left to right. 3. Because of the large size of the atoms at the bottom of a group compared to the atomic size at the top of the group, it would be easier (take less energy) to attach an electron to a metal with a large atomic radius than one with a small radius. It also means there is less affinity for an electron. One periodic trend, then, is that electron affinity decreases as we move down a group.
7 -7- Electron Affinity (cont.) For nonmetals, there is a natural tendency to want to gain electrons. Even for many metals, their electronic structure allows them to accept electrons even if it isn't their natural tendency. 1. If the E.A. is a large positive value, there is a strong attraction (affinity) for an electron. 2. If the E.A. is a small positive value, there is a weak attraction (affinity) for an electron, but it will accept an electron. 3. If the E.A. is a negative value, energy must be added to force an electron to be accepted by an atom. 4. A value of zero simply indicates that there is no tendency to gain the electron. EMPHASIZE: DIFFERENCE FROM BOOK - LOOKS AT THE D ENERGY INVOLVED IN THE PROCESS: If A + e > A - + energy (EXO) If A + e - + energy > A - (ENDO) IMPORTANT POINTS: 1. Even though the natural tendency for metals is to lose electrons, we must consider if there are orbitals free to accept electrons. Consider the sodium atom: configuration is [Ne] 3s' This means there is one electron in the single s orbital, but it also means there is room for one more electron. Na would be predicted to have a small positive value. 2. In contrast, the configuration of Mg is [Ne] 3s 2 This means there are two electrons in the single s orbital, and it also means there no room for any more electrons. To add an electron to magnesium means that there must be enough energy supplied to force the electron into one of the as-yet-unoccupied 3p orbitals. Note: AREAS OF NEGATIVE OR ZERO VALUES 1. at II-A 2. at II-B 3. at VIII All of these are due to the completion of a subshell. Note the low values at V-A and VII-B - half-filled subshells.
8 -8- IV. ELECTRONEGATIVITY - the tendency for an atom to attract electrons to itself when it is chemically combined with another element. - developed by Linus Pauling - based on the values of electron affinity and ionization energies - not experimental data like EA and IE. - high EN - nonmetals -low EN - metals -Note: there are no EN values for noble gases because these gases have no tendency to react! Trends: 1. lowest value - bottom left corner (Fr = 0.70) highest value - top right corner (F = 4.0) Actually the values were set with fluorine at 4.0 and all other values were calculated based on fluorine. 2. EN values tend to increase as you move across a period. 3. EN values tend to decrease as you move down a group or family. 4. Note the lack of change as you progress through the transition metals. V. DENSITY - remember that density is the ratio of mass to volume - mass increases regularly as you progress across a period - volume increases as you move down a group, but decreases you move across a period, but the amount of volume change is greater for the first four periods, but is not as great for the larger periods. - highest density then would be predicted to be in the bottom right corner, but because of the lanthanide contraction and the increased mass due to addition of the "f" subshell elements, the elements with the greatest density fit in the middle of the bottom, ie. Ir and Os have the greatest density. - review the significance of the lanthanide contraction Trends: V-formation - most dense elements are at the bottom of the V - lightest elements are hydrogen and helium VI. Melting / Boiling point Related to (1) size of atom; (2) electron arrangement and number of protons; (3) amount of attraction between atoms of the same element. Trends are not very regular: 1. Note that for halogens - as you move down the group, the elements move from the gaseous state to solids. 2. As you move across a period, you move from metals to nonmetals - MP tends to decrease 3. For the most part, MP and BP are related - high MP generally means high BP. 4. Elements of note: carbon - highest melting point of all elements tungsten - highest melting point of all metals Note: while tungsten also has a high BP, carbon's BP is not much higher than its MP.
9 -9- VII. Colors - transition metal salts are colored due to the movement of "d" electrons; all other salts are white - only two colored metals - Cu and Au - all others are silvery Note: all metals when ground into a fine powder appear black; only aluminum remains silvery. VIII. metallic / nonmetallic character - tendency of a metal is to lose electrons / nonmetals tend to gain electrons - would look at properties such as ionization energy and electron affinity francium should be the most metallic because of its low ionization energy, while fluorine would be the best nonmetal. Note: because of the EN values, all elements have a relative tendency for metallic and nonmetallic properties. In terms of electrical conductivity, silver is the best with copper coming in a close second. Gold is the third best electrical conductor. Notice that these three make up the coinage metals. Aluminum is 4th, by the way. IX. acid/base properties - metals tend to be basic while nonmetals tend to be acidic - especially true of oxides of the elements. - metalloids would then be amphoteric - able to react with either acids or bases. Remember that metallic oxides are classified as basic anhydrides; while nonmetallic oxides are acidic anhydrides. - for metals that have more than one charge, the higher oxidation state produces acidic properties Ex. chromium has 3 charges - Cr +6 is acidic; Cr +3 is amphoteric; Cr +2 is basic X. Oxidation Numbers (ionic charge) & Valence Electrons - Oxidation # s: the number of electrons that are typically gained/lost in a chemical bond in order to achieve an octet - Valence Electrons: # of electrons in outermost energy level - According to the Octet Rule, atoms are most stable (energetically) when 8 electrons occupy the valence shell - Ox. No and valence electrons for transition metals will not need to be memorized. Group # Ox. # # Valence e /
10 -10- An atom of Hydrogen consists of a single proton surrounded by an electron that resides in a spherical 1s orbital. Recall that orbitals represent probability distributions meaning that there is a high probability of finding this electron somewhere within this spherical region. This electron, being a negatively charged particle, is attracted to the positively charged proton. Now consider an atom of helium containing two protons (there are neutrons in the nucleus, too, but they are not pertinent to this topic) surrounded by two electrons both occupying the 1s orbital. In this case, and for all other many electron atoms, we need to consider not just the proton - electron attractions, but also the electron - electron repulsions. Because of this repulsion, each electron experiences a nuclear charge that is somewhat less than the actual nuclear charge. Essentially, one electron shields, or screens the other electron from the nucleus. The positive charge that an electron actually experiences is called the effective nuclear charge, Z eff, and Z eff is always somewhat less than the actual nuclear charge. Z eff = Z - S, where S is the shielding constant The two figures represent different instantaneous positions for the two electrons in an atom of helium. The electrons both occupy a 1s orbital meaning that the time-averaged positions of the electrons can be represented by a sphere. At any moment, however, we could envision the two electrons as being on opposite sides of the nucleus in which case they poorly shield each other from the positive charge. At a different moment, one electron may be between the nucleus and the other electron, in which case the electron farther from the nucleus is rather effectively shielded from the positive charge. Considering helium, you might initially think that Z eff would be one for each electron (i.e., each electron is attracted by two protons and shielded by one electron, 2-1 = 1), but Z eff is actually To determine Z eff, we do not subtract the number of shielding electrons from the actual nuclear charge (that would always result in a value of Z eff of one for any electron in any atom!), but rather we subtract the average amount of electron density that is between the electron we are concerned with and the nucleus. There are rules we can use to estimate the value of S and thus determine Z eff, but we do not need to go into such details for this course. Penetration An electron in an s orbital has a finite, albeit very small, probability of being located quite close to the nucleus. An electron in a p or d orbital on the other hand has a node (i.e., a region where there cannot be any electron density) at the nucleus. Comparing orbitals within the same shell, we say that the s orbital is more penetrating than the p or d orbitals, meaning that an electron in an s orbital has a greater chance of being located close to the nucleus than an electron in a p or d orbital. For this reason, electrons in an s orbital have a greater shielding power than electrons in a p or d orbital of that same shell. Also, because they are highly penetrating, electrons in s orbitals are less effectively shielded by electrons in other orbitals. For example, consider an atom of carbon whose electron configuration is 1s 2 2s 2 2p 2. The two electrons in the 1s orbital of C will do a better job of shielding the two electrons in the 2p orbitals than they will of shielding the two electrons in the 2s orbital. This means that for electrons in a particular shell, Z eff will be greater for s electrons than for p electrons. Similarly, Z eff is greater for p electrons than for d electrons. As a result, within a given shell of an atom, the s subshell is lower in energy than the p subshell which is in turn lower in energy than the d subshell.
MOSELEY and MODERN PERIODIC TABLE (designed by atomic numbers of elements)
MOSELEY and MODERN PERIODIC TABLE (designed by atomic numbers of elements) 1 PERIODS: Period number = Number of basic energy levels = The principal quantum number The horizontal lines in the periodic system
More informationElectron Configuration and Chemical Periodicity
Electron Configuration and Chemical Periodicity The Periodic Table Periodic law (Mendeleev, Meyer, 1870) periodic reoccurrence of similar physical and chemical properties of the elements arranged by increasing
More informationOrganizing the Periodic Table
Organizing the Periodic Table How did chemists begin to organize the known elements? Chemists used the properties of the elements to sort them into groups. The Organizers JW Dobereiner grouped the elements
More informationA few elements, including copper, silver, and gold, have been known for thousands of years
A few elements, including copper, silver, and gold, have been known for thousands of years There were only 13 elements identified by the year 1700. Chemists suspected that other elements existed. As chemists
More informationNihal İKİZOĞLU. MOSELEY and MODERN PERIODIC TABLE (designed by atomic numbers of elements) kimyaakademi.com 1
MOSELEY and MODERN PERIODIC TABLE (designed by atomic numbers of elements) kimyaakademi.com 1 PERIODS: Period number = Number of basic energy levels = The principal quantum number The horizontal lines
More informationChapter 8: Periodic Properties of the Elements
C h e m i s t r y 1 A : C h a p t e r 8 P a g e 1 Chapter 8: Periodic Properties of the Elements Homework: Read Chapter 8. Work out sample/practice exercises Check for the MasteringChemistry.com assignment
More informationThe Periodic Law Notes (Chapter 5)
The Periodic Law Notes (Chapter 5) I. History of the Periodic Table About 70 elements were known by 1850 (no noble gases) but there didn t appear to be a good way of arranging or relating them to study.
More informationNotes: Unit 6 Electron Configuration and the Periodic Table
Name KEY Block Notes: Unit 6 Electron Configuration and the Periodic Table In the 1790's Antoine Lavoisier compiled a list of the known elements at that time. There were only 23 elements. By the 1870's
More informationProfessor K. Section 8 Electron Configuration Periodic Table
Professor K Section 8 Electron Configuration Periodic Table Schrödinger Cannot be solved for multielectron atoms We must assume the orbitals are all hydrogen-like Differences In the H atom, all subshells
More informationPeriods: horizontal rows (# 1-7) 2. Periodicity the of the elements in the same group is explained by the arrangement of the around the nucleus.
The Modern Periodic Table 1. An arrangement of the elements in order of their numbers so that elements with properties fall in the same column (or group). Groups: vertical columns (#1-18) Periods: horizontal
More informationElectron Configurations and the Periodic Table
Electron Configurations and the Periodic Table The periodic table can be used as a guide for electron configurations. The period number is the value of n. Groups 1A and 2A have the s-orbital filled. Groups
More informationAccelerated Chemistry Study Guide The Periodic Table, Chapter 5
Accelerated Chemistry Study Guide The Periodic Table, Chapter 5 Terms, definitions, and people Dobereiner Newlands Mendeleev Moseley Periodic table Periodic Law group family period Page 1 of 38 alkali
More informationChapter 8. Periodic Properties of the Element
Chapter 8 Periodic Properties of the Element Mendeleev (1834 1907) Ordered elements by atomic mass Saw a repeating pattern of properties Periodic law when the elements are arranged in order of increasing
More informationUnit 2 - Electrons and Periodic Behavior
Unit 2 - Electrons and Periodic Behavior I. The Bohr Model of the Atom A. Electron Orbits, or Energy Levels 1. Electrons can circle the nucleus only in allowed paths or orbits 2. The energy of the electron
More informationDEVELOPMENT OF THE PERIODIC TABLE
DEVELOPMENT OF THE PERIODIC TABLE Prior to the 1700s, relatively few element were known, and consisted mostly of metals used for coinage, jewelry and weapons. From early 1700s to mid-1800s, chemists discovered
More informationChapter 6 The Periodic Table
Chapter 6 The Periodic Table Section 6.1 Organizing the Elements OBJECTIVES: Explain how elements are organized in a periodic table. Section 6.1 Organizing the Elements OBJECTIVES: Compare early and modern
More informationIntroduction period group
The Periodic Table Introduction The periodic table is made up of rows of elements and columns. An element is identified by its chemical symbol. The number above the symbol is the atomic number The number
More informationCHAPTER NOTES CHAPTER 14. Chemical Periodicity
Goals : To gain an understanding of : 1. Electron configurations 2. Periodicity. CHAPTER NOTES CHAPTER 14 Chemical Periodicity The periodic law states that when the elements are arranged according to increasing
More informationPage 1 of 9. Website: Mobile:
Question 1: Did Dobereiner s triads also exist in the columns of Newlands Octaves? Compare and find out. Only one triad of Dobereiner s triads exists in the columns of Newlands octaves. The triad formed
More informationChapter 7 Electron Configuration and the Periodic Table
Chapter 7 Electron Configuration and the Periodic Table Copyright McGraw-Hill 2009 1 7.1 Development of the Periodic Table 1864 - John Newlands - Law of Octaves- every 8 th element had similar properties
More informationLecture Presentation. Chapter 8. Periodic Properties of the Element. Sherril Soman Grand Valley State University Pearson Education, Inc.
Lecture Presentation Chapter 8 Periodic Properties of the Element Sherril Soman Grand Valley State University Nerve Transmission Movement of ions across cell membranes is the basis for the transmission
More informationCHAPTER 6 The Periodic Table
CHAPTER 6 The Periodic Table 6.1 Organizing the Elements Mendeleev: listed the elements in order of increasing atomic mass and in vertical columns according to their properties. Left blank spaces for undiscovered
More information2011 CHEM 120: CHEMICAL REACTIVITY
2011 CHEM 120: CHEMICAL REACTIVITY INORGANIC CHEMISTRY SECTION Lecturer: Dr. M.D. Bala Textbook by Petrucci, Harwood, Herring and Madura 15 Lectures (4/10-29/10) 3 Tutorials 1 Quiz 1 Take-home test https://chemintra.ukzn.ac.za/
More informationTest Review # 4. Chemistry: Form TR4-9A
Chemistry: Form TR4-9A REVIEW Name Date Period Test Review # 4 Location of electrons. Electrons are in regions of the atom known as orbitals, which are found in subdivisions of the principal energy levels
More informationChapter 7. Electron Configuration and the Periodic Table
Chapter 7 Electron Configuration and the Periodic Table Topics Development of the periodic table The modern periodic table Effective nuclear charge Periodic trends in properties of elements Electron configuration
More informationThe Periodic Table and Periodic Law
The Periodic Table and Periodic Law Periodic trends in the properties of atoms allow us to predict physical and chemical properties. Section 1: Development of the Modern Periodic Table Section 2: Classification
More informationPeriodic Relationships
Periodic Relationships 1 Tabulation of Elements Mendeleev (1869) Arranged by mass Tabulation by chem.& physical properties Predicted missing elements and properties 2 Modern Periodic Table Argon vs. potassium
More informationDiscovery of Elements. Dmitri Mendeleev Stanislao Canizzaro (1860) Modern Periodic Table. Henry Moseley. PT Background Information
Discovery of Elements Development of the Periodic Table Chapter 5 Honors Chemistry 412 At the end of the 1700 s, only 30 elements had been isolated Included most currency metals and some nonmetals New
More informationCHAPTER 6. Table & Periodic Law. John Newlands
CHAPTER 6 Table & Periodic Law 6.1 Developing a Periodic Table The periodic table was developed to show the properties of an element by simply looking at it's location. In 1860, chemists agreed on a way
More informationHonors Chemistry Unit 4 ( )
Honors Chemistry Unit 4 (2017-2018) Families (research and present) Metals/nonmetals Trends o Atomic radius o Electronegativity o Ionization energy o Metallic and nonmetallic character Review Ions Oxidation
More informationMade the FIRST periodic table
Made the FIRST periodic table 1869 Mendeleev organized the periodic table based on the similar properties and relativities of certain elements Later, Henri Moseley organized the elements by increasing
More informationChapter 6 - The Periodic Table and Periodic Law
Chapter 6 - The Periodic Table and Periodic Law Objectives: Identify different key features of the periodic table. Explain why elements in a group have similar properties. Relate the group and period trends
More informationPart I Assignment: Electron Configurations and the Periodic Table
Chapter 11 The Periodic Table Part I Assignment: Electron Configurations and the Periodic Table Use your periodic table and your new knowledge of how it works with electron configurations to write complete
More informationChapter 7 Electron Configuration and the Periodic Table
Chapter 7 Electron Configuration and the Periodic Table Copyright McGraw-Hill 2009 1 7.1 Development of the Periodic Table 1864 - John Newlands - Law of Octaves- every 8th element had similar properties
More informationTrends in the Periodic Table
Trends in the Periodic Table A trend is a predictable change in a particular direction. Example: There is a trend in the alkali metals to increase in reactivity as you move down a group. Atomic Radius
More informationChapter 5 - The Periodic Law
Chapter 5 - The Periodic Law 5-1 History of the Periodic Table I. Mendeleev's Periodic Table A. Organization 1. Vertical columns in atomic weight order a. Mendeleev made some exceptions to place elements
More information- Chapter 7 - Periodic Properties of the Elements
- Chapter 7 - Periodic Properties of the Elements Summary 7.1 Development of the periodic table 7.2 Effective nuclear charge 7.3 Size of atoms and ions 7.4 Ionization energy 7.5 Electron affinities 7.6
More informationNotes: Electrons and Periodic Table (text Ch. 4 & 5)
Name Per. Notes: Electrons and Periodic Table (text Ch. 4 & 5) NOTE: This set of class notes is not complete. We will be filling in information in class. If you are absent, it is your responsibility to
More informationTest Review # 5. Chemistry: Form TR5-8A. Average Atomic Mass. Subatomic particles.
Chemistry: Form TR5-8A REVIEW Name Date Period Test Review # 5 Subatomic particles. Type of Particle Location Mass Relative Mass Charge Proton Center 1.67 10-27 kg 1 +1 Electron Outside 9.11 10-31 kg 0-1
More informationChapter 4. Periodic Trends of the Elements. Chemistry: Atoms First Second Edition Julia Burdge & Jason Overby
Chemistry: Atoms First Second Edition Julia Burdge & Jason Overby Chapter 4 Periodic Trends of the Elements M. Stacey Thomson Pasco-Hernando State College Copyright (c) The McGraw-Hill Companies, Inc.
More informationPeriodic Table Workbook
Key Ideas: The placement or location of elements on the Periodic Table gives an indication of physical and chemical properties of that element. The elements on the Periodic Table are arranged in order
More informationThe Periodic Table. Chapter 5. I. History II. Organization III. Periodic Trends
The Periodic Table Chapter 5 I. History II. Organization III. Periodic Trends I. History P. 101-103 5a: The Periodic Table Material in chapter 5 is critical to understanding chapter 6! Early Organization
More informationTest Review # 4. Chemistry: Form TR4-5A 6 S S S
Chemistry: Form TR4-5A REVIEW Name Date Period Test Review # 4 Development of the Periodic Table. Dmitri Mendeleev (1869) prepared a card for each of the known elements listing the symbol, the atomic mass,
More informationPeriodic Relationships
Periodic Relationships 1 Tabulation of Elements Mendeleev (1869) Arranged by mass Tabulation by chem.& physical properties Predicted missing elements and properties 2 Modern Periodic Table Argon vs. potassium
More informationSearching for an Organizing Principle. Searching for an Organizing Principle. How did chemists begin to organize the known elements?
Searching for an Organizing Principle Searching for an Organizing Principle How did chemists begin to organize the known elements? Searching for an Organizing Principle A few elements, including copper,
More informationAtoimic Structure and the Periodic Table: Unit Objective Study Guide Part 2
Name Date Due Atoimic Structure and the Periodic Table: Unit Objective Study Guide Part 2 Directions: Write your answers to the following questions in the space provided. For problem solving, all of the
More informationFor the Periodic Table above indicate each of the following TRENDS: atomic size and ionic size. Na Na + F F - Ne < < < <
Chapter 6 Organizing the Elements THE PERIODIC TABLE AND PERIODIC LAW Periodic Table Summary Sheet For the Periodic Table above indicate each of the following TRENDS: atomic size and ionic size Na Na +
More informationnumber. Z eff = Z S S is called the screening constant which represents the portion of the nuclear EXTRA NOTES
EXTRA NOTES 1. Development of the Periodic Table The periodic table is the most significant tool that chemists use for organising and recalling chemical facts. Elements in the same column contain the same
More informationCh. 7- Periodic Properties of the Elements
Ch. 7- Periodic Properties of the Elements 7.1 Introduction A. The periodic nature of the periodic table arises from repeating patterns in the electron configurations of the elements. B. Elements in the
More informationChapter 7. Periodic Properties of the Elements. Lecture Outline
Chapter 7. Periodic Properties of the Elements Periodic Properties of the Elements 1 Lecture Outline 7.1 Development of the Periodic Table The periodic table is the most significant tool that chemists
More informationUNIT 5 THE PERIODIC TABLE
UNIT 5 THE PERIODIC TABLE THE PERIODIC TABLE EARLY ATTEMPTS OF CLASSIFICATION Many chemists started to organize and classify the elements according to their properties. In the 1790s, Antoine LaVoisier
More informationThe orbitals in an atom are arranged in shells and subshells. orbital 3s 3p 3d. Shell: all orbitals with the same value of n.
Shells and Subshells The orbitals in an atom are arranged in shells and subshells. n=3 orbital 3s 3p 3d Shell: all orbitals with the same value of n n=3 3s 3p 3d Subshell: all orbitals with the same value
More informationUnit 2 - Electrons and Periodic Behavior
Unit 2 - Electrons and Periodic Behavior Models of the Atom I. The Bohr Model of the Atom A. Electron Orbits, or Energy Levels 1. Electrons can circle the nucleus only in allowed paths or orbits 2. The
More informationELECTRON CONFIGURATION AND THE PERIODIC TABLE
ELECTRON CONFIGURATION AND THE PERIODIC TABLE The electrons in an atom fill from the lowest to the highest orbitals. The knowledge of the location of the orbitals on the periodic table can greatly help
More informationCh 8 Electron Configurations and Periodicity (Periodic table)
Ch 8 Electron Configurations and Periodicity (Periodic table) - An e 1 configuration is an atom s particular distribution of e 1 among the available subshells and orbitals. For example, the ground state
More information6.4 Electronic Structure of Atoms (Electron Configurations)
Chapter 6 Electronic Structure and Periodic Properties of Elements 317 Orbital n l m l degeneracy Radial nodes (no.) 4f 4 3 7 0 4p 4 1 3 2 7f 7 3 7 3 5d 5 2 5 2 Check Your Learning How many orbitals have
More informationSAMPLE PROBLEMS! 1. From which of the following is it easiest to remove an electron? a. Mg b. Na c. K d. Ca
SAMPLE PROBLEMS! 1. From which of the following is it easiest to remove an electron? a. Mg b. Na c. K d. Ca 2. Which of the following influenced your answer to number one the most? a. effective nuclear
More informationUnit 2 Periodic Table
2-1 Unit 2 Periodic Table At the end of this unit, you ll be able to Describe the origin of the periodic table State the modern periodic law Key the periodic table according to metals vs. nonmetals and
More informationPERIODIC CLASSIFICATION
5 PERIODIC CLASSIFICATION OF ELEMENTS TEXTBOOK, QUESTIONS AND THEIR ANSWERS Q.1. Do Dobereiner s triads also exist in the columns of Newland s octaves? Compare and find out. Ans. Triad of Li, Na and K
More informationChapter 7 Periodic Properties of the Elements
Chapter 7 Periodic Properties of the Elements Learning Outcomes: Explain the meaning of effective nuclear charge, Z eff, and how Z eff depends on nuclear charge and electron configuration. Predict the
More informationPeriodicity & Many-Electron Atoms
Chap. 8 ELECTRON CONFIGURAT N & CEMICAL PERIODICITY 8.1-8.2 Periodicity & Many-Electron Atoms Understand the correlation of electron configuration and the periodic character of atomic properties such as
More informationChapter 5 Notes Chemistry; The Periodic Law The Periodic Table The periodic table is used to organize the elements in a meaningful way.
Chapter 5 Notes Chemistry; The Periodic Law The Periodic Table The periodic table is used to organize the elements in a meaningful way. As a consequence of this organization, there are periodic properties
More informationThe History of the Modern Periodic Table. Modified from
The History of the Modern Periodic Table Modified from www.thecatalyst.com During the nineteenth century, chemists began to categorize the elements according to similarities in their physical and chemical
More informationX Chemistry: Periodic Classification of elements. Gist of Lesson for Quick Revision (By JSUNIL)
X Chemistry: Periodic Classification of elements Gist of Lesson for Quick Revision (By JSUNIL) 1. Periodic table: The table which classifies all the known elements on the basis of their properties in such
More informationUnit Five: The Periodic Table Ref:
Unit Five: The Periodic Table Ref: 10.11 11.2 11.4 History of P.T. Chlorine Bromine Iodine Dobrenier- (1829) Triads groups of three elements of similar chemical and physical properties. Cannizzarro (1860)
More information8.1 Early Periodic Tables CHAPTER 8. Modern Periodic Table. Mendeleev s 1871 Table
8.1 Early Periodic Tables CHAPTER 8 Periodic Relationships Among the Elements 1772: de Morveau table of chemically simple substances 1803: Dalton atomic theory, simple table of atomic masses 1817: Döbreiner's
More informationPeriodic Table. Engr. Yvonne Ligaya F. Musico 1
Periodic Table Engr. Yvonne Ligaya F. Musico 1 TOPIC Definition of Periodic Table Historical Development of the Periodic Table The Periodic Law and Organization of Elements in a Periodic Table Periodic
More informationTrends in Atomic Size. Atomic Radius-one half the distance between the nuclei of two atoms of the same element when the atoms are joined
Periodic trends Trends in Atomic Size Atomic Radius-one half the distance between the nuclei of two atoms of the same element when the atoms are joined Trends in Atomic Size Group Trend: Atomic radii of
More informationMODULE-21 TRENDS IN THE MODERN PERIODIC TABLE
MODULE-21 TRENDS IN THE MODERN PERIODIC TABLE Valency is defined as the number of electrons an atom requires to lose, gain, or share in order to complete its valence shell to attain the stable noble gas
More informationINTRODUCTORY CHEMISTRY Concepts and Critical Thinking Seventh Edition by Charles H. Corwin
Lecture INTRODUCTORY CHEMISTRY Concepts and Critical Thinking Seventh Edition by Charles H. Corwin The Periodic Table by Christopher G. Hamaker Illinois State University Classification of the Elements
More informationa) State modern periodic law. Name the scientist who stated the law.
INTEXT - QUESTION - 1 Question 1: a) State modern periodic law. Name the scientist who stated the law. b) What is a periodic table? How many groups and periods does modern periodic table have? Solution
More informationUnit Two Test Review. Click to get a new slide. Choose your answer, then click to see if you were correct.
Unit Two Test Review Click to get a new slide. Choose your answer, then click to see if you were correct. According to the law of definite proportions, any two samples of water, H2O, A. will be made up
More informationChapter 5 The Periodic Law
z Chapter 5 The Periodic Law z Section 5-1 History of the Periodic Table Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical
More informationThe Quantum Mechanical Model
Recall The Quantum Mechanical Model Quantum Numbers Four numbers, called quantum numbers, describe the characteristics of electrons and their orbitals Quantum Numbers Quantum Numbers The Case of Hydrogen
More informationMr. Dolgos Regents Chemistry PRACTICE PACKET. Unit 3: Periodic Table
*STUDENT* *STUDENT* Mr. Dolgos Regents Chemistry PRACTICE PACKET Unit 3: Periodic Table 2 3 It s Elemental DIRECTIONS: Use the reading below to answer the questions that follow. We all know by now that
More informationChapter 6: The Periodic Table
Chapter 6: The Periodic Table (Lecture Notes) Russian chemist Mendeleev proposed that properties of elements repeat at regular intervals when they are arranged in order of increasing atomic mass. He is
More informationPeriodic Classification of Elements
Periodic Classification of Elements Important Points: The first classification of elements is due to dobereiner in 1817. Dobereiner Triad Theory: "The atomic weight of the middle element is the arithmetic
More informationMendeleev s Table (1871) While it was the first periodic table, Mendeleev had very different elements, such as the very reactive potassium and the
Periodic Table Mendeleev s Table (1871) While it was the first periodic table, Mendeleev had very different elements, such as the very reactive potassium and the very stable copper, in the same family.
More informationTHE PERIODIC TABLE & PERIODIC LAW! Development of the Modern Periodic Table!
THE PERIODIC TABLE & PERIODIC LAW! Development of the Modern Periodic Table! Development of the Periodic Table! Main Idea: The periodic table evolved over time as scientists discovered more useful ways
More informationAtomic Structure. Ch 3 Prentice Hall
Atomic Structure Ch 3 Prentice Hall The Nuclear Atom By 1919 Rutherford concluded that the atom has a dense positive center called the nucleus containing what he called protons The electrons surround
More informationCHEM 1305: Introductory Chemistry
CHEM 1305: Introductory Chemistry The Periodic Table From Chapter 5 Textbook Introductory Chemistry: Concepts and Critical Thinking Seventh Edition by Charles H. Corwin Classification of Elements By 1870,
More informationOctober 05, Periodic_Trends_Presentation student notes.notebook. Periodic Trends. Periodic Trends: Effective Nuclear Charge
Periodic Trends: tomic Radius Ionization nergy lectronegativity Metallic haracter Ionic Radius Periodic Trends Five main trends in the periodic table will be discussed: The sizes of atoms Ionization energy
More informationPeriodic Trends. Elemental Properties and Patterns
Periodic Trends Elemental Properties and Patterns History of the Periodic Table 1871 Mendeleev arranged the elements according to: Increasing atomic mass Elements w/ similar properties were put in the
More informationPeriodic Properties (Booklet Solution)
Periodic Properties (Booklet Solution) Foundation Builders (Objective). (B) Law of triads states that in the set of three elements arranged in increasing order of atomic weight, having similar properties,
More informationSCH3U- R. H. KING ACADEMY ATOMIC STRUCTURE HANDOUT NAME:
Particle Theory of Matter Matter is anything that has and takes up. All matter is made up of very small. Each pure substance has its of particle, from the particles of other pure substances. Particles
More information- Some properties of elements can be related to their positions on the periodic table.
186 PERIODIC TRENDS - Some properties of elements can be related to their positions on the periodic table. ATOMIC RADIUS - The distance between the nucleus of the atoms and the outermost shell of the electron
More informationElectron Configuration and Periodic Trends - Chapter 5 section 3 Guided Notes
Electron Configuration and Periodic Trends - Chapter 5 section 3 Guided Notes There are several important atomic characteristics that show predictable that you should know. Atomic Radius The first and
More informationPERIODIC CLASSIFICATION OF ELEMENTS
PERIODIC CLASSIFICATION OF ELEMENTS Important terms & condition: Classification of elements: the arrangement of elements in such a manner that elements with similar properties are grouped together while
More information8.6,8.7 Periodic Properties of the Elements
Pre -AP Chemistry 8.6,8.7 Periodic Properties of the Elements READ p. 305 315, 294-296 Practice Problems Pg 315 -Exercise 8.9 Pg 318-321 #36, 55, 64, 66, 67, 69, 72, 80 Periodic Trends are predictable
More informationThe Atom & Periodic Table. Unit 2 Topics 4-6
The Atom & Periodic Table Unit 2 Topics 4-6 Electrons in Atoms Topic 4 Describe Bohr s model of the atom. Sketch it! Bohr - A review electrons exist in orbits around the nucleus. Bohr - IB Information
More informationPowerPoint to accompany. Chapter 6. Periodic Properties of the Elements
PowerPoint to accompany Chapter 6 Periodic Properties of the Elements Development of the Periodic Table Elements in the same group generally have similar chemical properties. Properties are not identical,
More informationThe Periodic Table and Periodic Trends
The Periodic Table and Periodic Trends The properties of the elements exhibit trends and these trends can be predicted with the help of the periodic table. They can also be explained and understood by
More informationAdvanced Chemistry. Mrs. Klingaman. Chapter 5: Name:
Advanced Chemistry Mrs. Klingaman Chapter 5: The Periodic Law Name: _ Mods: Chapter 5: The Periodic Law Reading Guide 5.1 History of the Periodic Table (pgs. 125-129) 1) What did Dimitri Mendeleev notice
More informationUnit 7 Study Guide: Name: KEY Atomic Concepts & Periodic Table
Unit 7 Study Guide: Name: KEY Atomic Concepts & Periodic Table Focus Questions for the unit... How has the modern view of the atom changed over time? How does a chemist use symbols and notation to communicate
More informationChapter 7 The Structure of Atoms and Periodic Trends
Chapter 7 The Structure of Atoms and Periodic Trends Jeffrey Mack California State University, Sacramento Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS
More informationClass: XI Chapter 3: Classification of Elements Chapter Notes. Top Concepts
Class: XI Chapter 3: Classification of Elements Chapter Notes Top Concepts 1. Johann Dobereiner classified elements in group of three elements called triads. 2. In Dobereiner s triad the atomic weight
More informationRegan & Johnston Chemistry Unit 3 Exam: The Periodic Table Class Period
Regan & Johnston Name Chemistry Unit 3 Exam: The Periodic Table Class Period 1. An atom of which element has the largest atomic radius? (1) Si (2) Fe (3) Zn (4) Mg 2. Which characteristics both generally
More informationDeveloping the Periodic Table
Developing the Periodic Table Early Element Classification Mendeleev s First Periodic Table Mendeleev s First Periodic Table Mendeleev s Periodic Table Arranged by increasing atomic mass Some elements
More informationChapter 7. Generally, the electronic structure of atoms correlates w. the prop. of the elements
Chapter 7 Periodic Properties of the Elements I) Development of the P.T. Generally, the electronic structure of atoms correlates w. the prop. of the elements - reflected by the arrangement of the elements
More informationUnit III. Chemical Periodicity
Unit III. Chemical Periodicity History of Periodic Table Modern Periodic Table Periodic Trends Links Return to Notes PDF Format READING ASSIGNMENT 1: Read Ch. 14.1, p 391-396. Answer questions 1-5 p. 396
More information