Untitled Flashcards Set

1. Modern Atomic Theory

Periodic Table of Elements

• Elements were first arranged by increasing atomic number

• Periodic table was arraigned before electron configurations of elements was known 

• Properties of elements are periodic

• Dmitri Mendeleev (Russia)

• Lothar Meyer (Germany)

Features of Periodic Table 

1. Periods

2. Rows

3. Metals

4. Nonmetals

5. Metalloids 

6. Hallogens

   1. Group 7a

7. Noble Gases 

   1. Group 8a

8. Alkali metals 

   1. Group 1A

Metals

• Physical Properties

  • Mostly solid states at standard temperature and pressure

  • Conduct electricity

  • Ductile and malleable

• Chemical Properties

  • Form cations 

  • Reducing agents

  • Basic oxides or amphoteric

Non-Metals 

• Physical Properties

  • Vary in state at standard temperature 

  • Do not conduct electricity (except graphite form in Carbon)

  • Varying properties

• Chemical properties

  • Forming anions

  • Oxidizing agents

  • Acidic oxides

Metalloids 

• “Intermediates of Metals and Nonmetals”

• May conduct electricity like metals

• May resemble nonmetals

Structure of Atoms

• Subatomic Particles 

  • Contains Nucleus, Electron, and Protons

  • Electron cloud; somewhere there is the electron

  • Neutron: no charge

  • Electron: negative

  • Proton: positive charge

• Atomic Composition: The atom is mostly empty space

  • The protons and neutrons in the nucleus

  • The number of electrons is equal to the number of protons (electronically neutral)

  • Electrons in space around nucleus (electron cloud)

  • If we’re dealing with atomic mass, the electron has a small effect on it, since it has almost negligible mass

  • Summary:

    • Atoms are neutral: number of protons and electrons must be equal

  • Atomic mass = p + n

    • Don’t have to worry about mass of e since they have such a small mass

    • Number of neutrons are determined from atomic mass

• Atomic Number, Z

  • All atoms of the same element have the same number of protons in the nucleus, Z

• Atomic Weight

  • The atomic mass of one atom of an element is relative to one atom of another

  • Standard used is C-12 isotope

  • Unit:

    • Atomic mass unit (u)

    • 1 u = 1.66054 x 10^-12 g

• Mass number 

  • A = mass number = number of protons + number of neutrons

• Note: the subscript Z is optional because the elemental symbol tells us what the atomic number must be

• Can also be denoted as X-A i.e. C-12

• Summary of Mass Number

  • C-12 atom has 16 protons and 6 neutrons

  • Mass number (A) = # of protons + # of neutrons

  • Carbon - 12  has a mass # of 12 u

  • Carbon -13 has a mass # of 13 u (6 +7) and so on.

  • Elements with different number of neutrons are called “Isotopes”

Answer for Practice Problem

Isotopes

• Existence of isotopes means all atoms of an element are not exactly the same

• Existence of isotopes means all atoms of an element are not exactly the same

• Isotopes have same atomic number (Z) but different total  number of nucleons (A)

Isotopes are measured through

• The mass spectrometer gives information on the mass and relative abundance of each element’s isotopes

• Each isotope is represented by a Relative Abundance/Percent Abundance

Atomic Weight (Mass) 

• The atomic masses on the periodic table are “weighted averages” of the all of an element’s individual isotope masses.

• All of the given isotopes should be added to each other

• Atomic mass is 4 sig figs since the percent is XX.XX%

Chemical Nomenclature

• Important when giving entire compounds being spelled 

Types of Compounds

• Ionic compounds

  • Formed with the reaction between metals & nonmetals

• Consists of Ions (atomic or groups of atoms) that bear a positive or negative electric charge

• These are generally referred to as salts

Ions

• Are atoms or groups of atoms with a formal positive or negative charge

• Removing electrons produces a cation (+)

• Adding electrons produces Anion (-)

Predicting Ion Charges 

• In general 

  • Metals lose electrons forming cations

  • Nonmetals gain electrons forming anions

• By losing or gaining electrons, an atom has same number of electrons as the nearest noble gas atom

Stock System for Nomenclature

• Cation name of an element followed by ion

  • Ex. Sodium ion

• Anion name of the element with an -ide ending

  • Ex Fluoride, Chloride, etc

• Transition metals have multiple charge states

  • Element + (charge in roman numeral) + ion

    • Ex. Cu⁺ → Copper (I) ion

    • Ex. Cu²⁺ → Copper (II) ion

Polyatomic ions 

• A special class of ions where a group of atoms tend to stay together - an ion that constraints atoms covalently bound together

• Polyatomic anions are groups of atoms (molecules) with a net charge 

Forming and Naming Ionic Compounds 

Cation name + anion name

• For main group metals, no need to indicate charge 

• For transition metals, need to indicate charge

• Ionic compounds are electrically neutral (no net charge)

• Use the crossover rule (apples to mono and polyatomic ions)

Molecular Formulas

• Molecular formula: #’s & types of each atom

• Condensed formula: indicates how certain atoms are grouped together

• Structural formula: Shows the connection (bonds)

• Molecular model: Give 3-D perspective

Naming Molecular Compounds

When nonmetals combine they form molecules. They may do so in multiple forms 

Ex. CO → carbon monoxide

CO₂ → carbon dioxide

Because of this we need to specify the number of each atom by way of prefix 

Examples:

Development of Atomic Theory

2. Electron Configuration

Quantum Mechanical Model: Bohr, Schrodinger, and Heisenberg

• The line spectra experiment proved to Bohr that an atom only has certain discrete energy levels, and any movement from these levels is due to energy adsorption and emission 

• Schrodinger formulated Schrodinger’s Equation, wave functions that act as mathematical models which postulate the existence of the electron at different energy levels.

  • This also is the basis for the Quantum Mechanical Model of the Atom

• Heisenberg created the Uncertainty principle, wherein we cannot know the exact location of the electron at any given moment. But by using the Schrodinger’s Equation we can know where it may possibly be

  • Where you can only know the energy and the position 

  • We can postulate where the position of an electron is

The atom - composed of Nucleus (protons and neutrons) as its center and an electron cloud

The analogy of the electron probability density

• Take an apple tree 

• Apples may be found near the trunk of the tree

• While it can be found around the trunk, the farther you go the less apples there are

• We can expect an apple nearer from the trunk

• “The apple never falls far from the tree”

The electron probability density 

• Schrodinger’s equation is a wave function, which gave a mathematical description of the probability of the location of the electron

  • Malalaman natin kung asan ba ang electron sa atom

• To describe the electron of an element, three quantum numbers are needed. The set of these numbers is called atomic orbital

Quantum Numbers: Numbers used to describe the electron of the element

• First three are the address and the last one is the property of the 

1. Principal quantum number (n): Described as the shell

n = 1,2,3,4,…..n

2. Angular quantum number (l): Described as the subshells

l = 0,1,2,3, n-1

3. Magnetic Quantum Number (ml): Described as the orbitals

ml = 1,(-I,0,I),.....2l+1

4. Electron-spin quantum number (m_s): Described as the 

ms = +½ or -½

Principal Quantum Number (n)

• The energy level or shells of the atom

• The period in the periodic table

• Indicates relative size of the orbital

• Relative distance from nucleus of the peak in the radial probability distribution plot

• N increases, orbital becomes larger, electron becomes farther

• A positive integer (n=1,2,3,...)

• If papalayo ka ng paplayo sa nucleus, the probability of the electron gets lower

• Lower probability = Larger orbital 

Angular Quantum number (l)

• Angular quantum number (l)

• l is an integer from 0 to n-1

  • l = 0,1,2,3,...n-1

• Gives the characteristic shape of the orbital

• The block in the periodic table

  • Look at the graphic below

• The number of subshells given by the value of l

Types of Orbitals

• Spherical (s orbital)

  • l=0

  • Sphere = globe like

• Dumbbell-like (p orbital)

  • l=1

  • Shape is like a dumbbell or 2 circular shapes that goes on both sides

• Cloverleaf & Donut (d orbital)

  • l=2

• Tetrahedral (f orbital)

  • l=3

• The higher the value of the orbital the more intricate the shape of the orbital

• The higher the l value the harder it is for the electron to find 

• In relation to housing electrons 

  • “s” orbitals can hold 2 electrons

  • “p” orbitals can hold up to 6 electrons

  •  “d” orbitals can hold up to 10 electrons

  • “f” orbitals can hold up to 14 electrons

Magnetic Quantum Number (m_l)

• m_l is an integer

  • -l,0,l; (-l,0,l)

  • m_l = 2l+1

• Prescribes the 3D orientation of the orbital in space around nucleus based on the value of l

• Corresponds to the different orientations of your orbitals

Electron-spin quantum number (m_s)

• Property of the electron, not the orbital

• m_s is a number that describes the spin of the electron

  • Whether it’s going up or down

• It may have either the value of +½ and -½ 

• What is the use for this quantum number?

  • One orbital can house 2 electrons

  • But in cases where we need to house 2 electron, it gives us a way to distinguish both electrons

Hierarchy of QN for Atomic Orbital

SUMMARY of Quantum Numbers

QUESTION: How did we get the sublevel name?

We get the name by finding out the sublevel delegation, the sublevel delegation is show below

Electron Configuration

• Shows the distribution of electrons within the levels and sublevels of its atoms

Principles for Electron Distribution

1. Pauli’s Exclusion Principle

No two electrons in the same atom can have the same four quantum numbers

Ex.

• Basta dapat iba siya

2. Aufbau Principle

Fill every orbital of the lowest sublevel first before moving up

• Filling electrons are assigned to subshells in order of increase “n + 1” value

• For two subshells with the same value of “n+l”, electrons are assigned first to the subshell of lower n

Ex: 

4s: n + l = 4 + 0 =4

3d: n + l = 3 +2 = 5

This is why 4s goes first since it has less energy than 3d, but 4s is always going to be in the outer electron.

3. Hund’s Rule

Fill every orbital in the same sublevel first before any pairing occurs.

Type of Electrons 

• Inner (core) electrons

  • Electrons an atom has in common with the previous noble gases 

  • [Ne], [Ar], [Kr]

• Outer electrons

  • Those in the highest energy level (highest n value); farthest away from the nucleus 

• Valence electrons

  • Involves in forming compounds 

  • For main-group elements, the valence electrons are the outer electrons

  • For transition elements, aside from ns and np electrons, elements Z=26 to Z=30 can use their d electrons for bonding.

Electron Configuration

Shows the distribution of electrons within the levels and sublevels of its atoms 

1. Full Electron Configuration: Ca (Z=20)

2. Core Notation

• Write the nearest noble gas, and then writing the sublevels afterwards

3. Orbital diagram

Period 1 and Period 2

Ions and Isoelectronic

• Ions are elements who either lose or gain electrons. But again, why do they want to lose/gain their electrons?

• Because they want to be noble gases by being isoelectronic with them (same electron configuration)

• Isoelectronic - the same electron configuration as a noble gas

• Example:

Period 4 Transition Series

Filling up electrons for the transition elements series are a bit tricky

Three factors:

• Effects of shielding and penetration on sublevel energy

• Filling the 4s and 3d sublevels

• Stability of half-filled and filled sublevels

For + Ions 

• You remove first the highest level given in for example: 

  • Fe = [Ar] 4s² 3d⁶

  • But if we turn it into Fe³⁺

  • It will become Fe = [Ar] 3d⁵ since it lost 3 electrons and the first one to lose the electrons are from the 4s² 

PERIODIC TRENDS 

Periodicity 

• Periodic representation of the properties of elements

• Observable trends between neighboring elements in the periodic table

Coulomb’s Law 

• Explains the electrostatic interactions between charges

  • Like charges repel, unlike charges attract

• Effect of Nucleus - electron attraction

  • Pag meron ka friend sa UP or ADMU san ka mas close

• Effects of Electron - repulsion

  • Shielding - 

    • There are repulsive forces happening to two electrons,

Electrostatic interaction determine sublevel energies due to:

Your orbital type has an effect to the location of your relectron 

• Penetration: the reason why 4s is nauuna kaysa sa 3d

• Metalloids can have both properties of metals and nonmetals 

Electron Affinity

• Energy change accompanying the addition of 1 mol of electron

Trends in Electron Affinity

List of questions 

7. In general, ionization energy increases from left to right across a given period. Aluminum, however, has a lower ionization energy than magnesium. Explain in no more than 4 sentences. [3 pts]

Aluminum has a lower ionization energy than Magnesium because Magnesium has a full 3s electron shell, making it harder to remove or add ions to it. Aluminum on the other hand has one electron in its 3p orbital, making it easier to add or remove electrons in its valence shell, because the 3p Aluminum electron is farther from the nucleus and experiences more shielding. This in turn makes removing the electron to require less energy.

8. Explain the effect of the interaction between electrons and nucleus to effective nuclear charge and consequently, to the atomic radius trend. Explain in no more than 5 sentences [5 pts].

Electrons are attracted to the positively charged nucleus, but inner electrons shield outer electrons from the full nuclear charge, reducing the effective nuclear charge experienced by the valence electrons. Once you go up the periods, Zeff increases because protons are added without additional shielding, pulling electrons closer and decreasing atomic radius. Moving down a group, additional electron shells increase shielding, reducing Zeff and causing an increase in the atomic radius. Higher Zeff results in a stronger attraction between the nucleus and electrons, leading to a smaller atomic radius, while lower Zeff leads to a larger atomic radius. Thus, atomic radius decreases across a period and increases down a group in the periodic table.