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Periodic Trends

This text explains the trends in atomic and ionic radius, electronegativity, ionization energy, and density of elements based on the periodic table. It discusses why these properties increase or decrease across periods and groups.

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Periodic Trends

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  1. Periodic Trends SC1. Obtain, evaluate, and communicate information about the use of the modern atomic theory and periodic law to explain the characteristics of atoms and elements. f. Use the periodic table as a model to predict the relative properties of elements based on the patterns of electrons in the outermost energy level of atoms (i.e. including atomic radii, ionization energy, and electronegativity).

  2. ATOMIC RADIUS • Atomic radius is defined as one-half of the distance between covalently bonded nuclei. • Usually measured in picometers (1 pm = 10-12 m) or angstroms (1 Å = 10-10 m).

  3. ATOMIC RADIUS

  4. ATOMIC RADIUS

  5. ATOMIC RADIUS Why? Each step adds a proton and an electron, making the nucleus more positive and the electron cloud more negative. The increased attraction pulls the cloud in. Why? With each new period, we add an entirely new energy level, making the atoms larger with each step. INCREASES INCREASES

  6. IONIC RADIUS • Ionic radius is defined as the radius of an ion. • Cations are smaller than the neutral atom • Why? The entire outermost energy level is removed. • Anions are larger than the neutral atom • Why?Repulsion resulting from the additional electron(s) enlarges the electron cloud.

  7. IONIC RADIUS

  8. IONIC RADIUS

  9. IONIC RADIUS Why? Each step increases the effective nuclear charge, making the nucleus more positive, causing the remaining electrons to be pulled closer. Why? With each new period, we add an entirely new energy level, making the atoms larger with each step. INCREASES INCREASES

  10. ELECTRONEGATIVITY • Electronegativity is the tendency of an atom or attract electrons in a bond and, thus, the tendency to form negative ions. • It is an arbitrary scale that ranges from 0 to 4. • Measured in Paulings • Generally, metals form cations and have low electronegativities. • Nonmetals typically form anions and have high electronegativities (F has the highest value of 4).

  11. ELECTRONEGATIVITY

  12. ELECTRONEGATIVITY Why? The effective nuclear charge increases, causing electrons to be attracted to that atom. Why? As more energy levels are added, the inner (non-valence) electrons shield the outer (valence) electrons from the attraction of the nucleus. INCREASES INCREASES

  13. IONIZATION ENERGY • Ionization energy is the energy required to remove a valence electron from an atom. • Measured in electron volts (eV) or kilojoules (kJ) • It requires more energy to remove each successive electron. • Generally, metals have low IEs and nonmetals have high IEs

  14. IONIZATION ENERGY

  15. IONIZATION ENERGY Why? As the nucleus becomes more positive, electrons are pulled closer, making it harder to remove the electrons from the atom. Why? With each new period, we add an entirely new energy level, making the atoms larger with each step and increasing the shielding effect. This reduces the nuclear pull on the electrons. INCREASES INCREASES

  16. DENSITY • Density is defined as the ratio of mass to volume. • Measured in g/cm3 or kg/L • Solids (specifically metals) have high densities • Gases have low densities

  17. DENSITY

  18. DENSITY Why?(A) Gases have lower densities than solids. (B) Among solids, transition metal atoms pack more closely as the atoms themselves grow smaller (atomic radius trend). Why? Mass increases greatly due to increased size of the nucleus, but the volume (electron cloud) does not change as much. INCREASES

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